A  COURSE  IN 
GENERAL  CHEMISTRY 


BY 

WILLIAM  McPHERSON 
\\ 

AND 

WILLIAM  EDWARDS  HENDERSON 

PROFESSORS   OF   CHEMISTRY,  OHIO  STATE   UNIVERSITY 


GINN  AND  COMPANY 

BOSTON  •  NEW  YORK  •  CHICAGO  •  LONDON 


COPYRIGHT,  1913,  BY 
WILLIAM  McPHERSON  AND  WILLIAM  E.  HENDERSON 


ALL   RIGHTS   RESERVED 
514.3 


gfte   fltftenatum    -Qrtfg 

GINN  AND  COMPANY  •  PRO- 
PRIETORS •  BOSTON  •  U.S.A. 


PEEFACE 

In  preparing  this  textbook  in  general  chemistry  the  authors  have 
been  guided  by  a  few  simple  principles.  They  have  endeavored  to 
write  a  book  scientific  in  spirit  and  at  the  same  time  thoroughly  teach- 
able. It  may  be  said  at  once  that  the  text  presents  few  novelties,  either 
in  arrangement  or  in  method.  While  individual  teachers  can  succeed 
better  by  adopting  unusual  and  original  modes  of  presentation,  we 
appear  to  have  reached  a  more  or  less  standard  method,  which  is,  on 
the  whole,  best  adapted  to  the  greatest  number. 

The  most  noticeable  departure  from  this  arrangement  in  the  pres- 
ent text  is  the  postponement  of  the  halogen  elements  to  a  relatively 
late  chapter.  Several  considerations  of  pedagogy  have  prompted  this 
change :  (1)  The  student  needs  some  training  in  experimental  manip- 
ulation before  attempting  laboratory  exercises  with  these  active  ele- 
ments. (2)  The  action  of  these  elements  upon  the  alkalies,  and  the 
conduct  of  their  oxygen  acids,  involve  many  difficulties,  and  the 
attempt  to  meet  these  by  presenting  the  chemistry  of  these  elements 
in  several  widely  separated  chapters  is  bewildering  to  the  student. 
(3)  It  is  unsatisfactory  to  develop  the  conception  of  an  acid  by  means 
of  the  example  of  the  halogen  hydrides,  since  they  are  far  from  being 
typical  acids.  In  the  experience  of  the  authors  the  order  followed  in 
this  text  is  much  to  be  preferred. 

In  the  main  the  elements  have  been  grouped  in  description  in  ac- 
cordance with  their  positions  in  the  periodic  classification,  but  the 
authors  have  not  hesitated  to  abandon  this  order  when  a  more  natural 
one  suggested  itself.  Thus,  with  the  metals  the  arrangement  chosen 
presents  the  strongly  electropositive  elements  of  unchanging  valence 
first.  These  are  followed  by  those  which  present  the  complications  of 
changeable  valence,  sulfide  metallurgy,  and  amphoteric  characteristics. 

Every  teacher  desires  to  impress  his  students  with  the  truth  that 
science  rests  fundamentally  upon  experiment  and  is  therefore  subject 
to  the  uncertainties  of  experimental  error.  The  authors  believe  that 
this  end  is  best  attained  by  the  free  introduction  of  historical  matter. 

iii 


288fifl.ni 


iv  GENERAL  CHEMISTRY 

This  gives  to  abstractions  a  personal  touch  which  makes  them  seem 
human  —  more  a  product  of  experience  and  imagination  and  less  of  the 
nature  of  a  revelation.  While  the  adoption  of  this  method  has  added 
materially  to  the  size  of  the  book,  it  is  believed  that  it  has  increased 
its  interest  and  at  the  same  time  kept  it  faithful  to  the  spirit  of  science. 

The  treatment  of  theory  in  any  text  will  always  leave  many  teachers 
dissatisfied.  Each  has  his  own  point  of  view,  and  the  authors  have 
simply  been  loyal  to  their  own.  They  have  had  no  partisan  feeling 
for  any  particular  theory,  but,  on  the  contrary,  have  consistently  tried 
to  impress  the  student  with  the  idea  that  theories  are  largely  a  matter 
of  convenience  and  are  not  to  be  regarded  as  final. 

A  rather  unusual  amount  of  material  has  been  introduced  which 
might  fairly  be  left  to  the  course  in  physics.  With  the  elective  system 
so  widely  adopted  in  American  colleges  it  can  no  longer  be  assumed 
that  a  student  has  had  a  course  in  physics  before  entering  upon  the 
study,  of  chemistry.  Moreover,  it  is  a  rather  common  experience  to 
find  that  the  knowledge  which  a  student  has  acquired  in  one  course  is 
with  some  difficulty  called  to  his  assistance  in  another. 

The  extent  to  which  the  elements  of  organic  chemistry  should  be 
incorporated  into  the  general  course  must  always  be  decided  by  local 
conditions.  The  present  text  includes  a  chapter  on  the  compounds  of 
carbon,  which  the  authors  feel  is  in  good  proportion  to  the  other 
chapters.  Most  of  this  may  be  omitted  if  the  requirements  of  a  course 
render  this  desirable. 

In  the  matter  of  spelling  the  various  chemical  terms,  the  authors 
have  been  conservatively  modern.  The  shorter  method  for  spelling 
sulfur  is  now  so  generally  used  that  there  seems  to  be  no  good  reason 
opposed  to  its  universal  adoption. 

The  growing  importance  in  the  industries  of  some  of  the  rarer 
metals  renders  a  more  intimate  knowledge  of  their  properties  desirable. 
Accordingly,  the  discussion  of  these  metals  is  somewhat  more  extended 
than  is  usual  in  a  text  of  this  scope. 

The  authors  have  frankly  adopted  all  the  devices  of  paragraphing 
and  typography  which  would  make  the  logical  argument  and  the 
coordination  of  matter  stand  out  prominently.  They  have  also  en- 
deavored to  have  the  drawings  true  to  scale  and  representing  ordinary 
laboratory  apparatus.  They  were  made  by  Mr.  Cree  Sheets  under 
the  direction  of  our  colleague,  Professor  Thomas  E.  French,  of  the 
department  of  engineering  drawing. 


PREFACE  v 

No  effort  has  been  spared  to  have  the  statements  concerning  indus- 
trial processes  accord  with  present  usage,  and  we  are  indebted  to 
our  colleague,  Professor  James  R.  Withrow,  as  well  as  to  the  managers 
of  many  industrial  concerns,  for  much  information  along  these  lines. 
Other  colleagues  have  given  us  the  benefit  of  expert  knowledge  in 
special  topics,  among  whom  we  desire  to  name  Professor  Edward 
Orton,  Jr.,  department  of  ceramics ;  Professor  Edward  E.  Somermeier 
and  Professor  Dana  J.  Demorest,  department  of  metallurgy ;  and 
Professor  Charles  B.  Morrey,  department  of  bacteriology.  Finally,  we 
owe  much  to  the  experience  and  kindly  interest  of  the  associates  in  our 
own  department,  Professors  Charles  W.  Foulk  and  William  L.  Evans, 
and  Dr.  David  R.  Kellogg,  who  has  corrected  all  the  proof  sheets. 
To  all  those  whom  we  have  mentioned,  and  to  many  other  friends 
as  well,  we  here  express  our  deep  sense  of  gratitude.  At  the  same 
time  we  wish  to  emphasize  the  fact  that  the  authors  alone  are  respon- 
sible for  any  errors  which  may  be  found  in  the  text. 

THE  AUTHORS 
THE  OHIO  STATE  UNIVERSITY 


CONTENTS 


CHAPTER                                                                                                                                                  .  PAGE 

I.   MATTER  AND  ENERGY 1 

II.   OXYGEN 16 

III.  HYDROGEN 35 

IV.  PROPERTIES  OF  GASES 45 

V.   WATER 54 

VI.   THE  THREE  STATES  OF  MATTER 73 

VII.    THE  LAWS  OF  CHEMICAL  COMBINATION  :  THE  ATOMIC  THEORY  83 

VIII.    EQUATIONS  AND  CALCULATIONS 98 

IX.   NITROGEN  AND  THE  RARE  ELEMENTS:  HELIUM,  NEON,  ARGON, 

KRYPTON,  XENON 104 

X.    THE  ATMOSPHERE 114 

XL    SOLUTIONS  ....  122 

XII.     ION1ZATION  AND  ELECTROLYSIS 137 

XIII.  NEUTRALIZATION 148 

XIV.  VALENCE  AND  STRUCTURAL  FORMULAS 161 

.  XV.    COMPOUNDS  OF  NITROGEN 167 

XVI.    EQUILIBRIUM 191 

XVII.    SULFUR;  SELENIUM;  TELLURIUM 203 

XVIII.   CLASSIFICATION  OF  THE  ELEMENTS 233 

XIX.    THE  CHLORINE  FAMILY 243 

XX.    THE  OXYGEN  COMPOUNDS  OF  THE  HALOGENS 268 

XXL    CARBON  AND  ITS  COMPOUNDS 275 

VKXII.    MOLECULAR  WEIGHTS 309 

XXIII.  FLAMES;  FUEL  GASES;  EXPLOSIONS 322 

XXIV.  THERMOCHEMISTRY 332 

XXV.   THE  SILICON  FAMILY  AND  BORON 341 

XXVI.    THE  PHOSPHORUS  FAMILY     . 354 

XXVII.    THE  HYDROXIDES  AND  THEIR  REACTIONS 377 

XXVIII.    THE  METALS 382 

XXIX.    THE  ALKALI  METALS 391 

vii 


viii  GENERAL  CHEMISTRY 

CHAPTER  PAGE 

XXX.  THE  ALKALINE  EARTH  METALS 413 

XXXI.  THE  MAGNESIUM  FAMILY .     .  429 

XXXII.  THE  ALUMINIUM  GROUP 439 

XXXIII.  THE  SILICATE  INDUSTRIES 453 

XXXIV.  THE  IRON  FAMILY 458 

XXXV.  COPPER  ;  MERCURY  ;  SILVER 477 

XXXVI.    TIN  AND  LEAD 500 

XXXVII.   MANGANESE  AND  CHROMIUM 514 

XXXVIII.    THE  VANADIUM  AND  MOLYBDENUM  FAMILIES      .....  525 

XXXIX.    GOLD  AND  THE  PLATINUM  FAMILY 536 

INDEX ,-v.     543 

APPENDIX  A Facing  back  cover 

APPENDIX  B  liiside  back  cover 


GENERAL  CHEMISTRY 

CHAPTER  I 

MATTER  AND  ENERGY 

Introduction.  We  early  learn  in  common  experience  that  all  material 
things  with  which  we  are  acquainted  are  constantly  undergoing  changes 
of  some  sort.  Each  succeeding  season  profoundly  alters  the  whole 
aspect  of  nature  around  us.  Even  those  things  which  we  consider  to 
be  the  most  durable  in  time  show  evidences  of  change.  The  metals 
rust  away ;  the  solid  rocks  weather  and  disintegrate ;  the  most  im- 
posing buildings  crumble  into  ruins.  The  variety  of  such  changes  in 
nature  is  almost  without  limit,  yet  it  is  possible  to 'resolve  them  all 
into  two  fundamental  factors  —  the  matter  which  undergoes  change, 
and  the  energy  which  occasions  the  change. 

Many  of  these  changes  result  in  the  formation  of  substances  which 
are  at  once  recognized  to  be  different  from  the  original  materials. 
When  iron  rusts  it  is  evident  that  the  rust  is  quite  different  from  the 
iron.  When  a  lump  of  coal  burns,  the  invisible  gases  and  handful  of 
ashes  into  which  it  is  converted  bear  no  resemblance  to  the  coal. 
These  examples  suggest  many  questions  as  to  the  real  nature  of 
these  changes,  why  they  occur,  and  in  what  respects  the  products 
differ  from  the  original  materials.  The  science  of  chemistry  has  for 
its  object  the  accurate  investigation  of  all  such  changes,  the  causes 
which  occasion  them,  the  energy  changes  which  accompany  them,  and 
the  laws  in  accordance  with  which  they  take  place.  Chemistry  is 
therefore  very  intimately  concerned  with  both  matter  and  energy, 
and  it  will  be  of  much  advantage  if  at  the  outset  we  get  clearly 
before  us  some  fundamental  conceptions  in  regard  to  them.  We 
shall  then  be  in  a  better  position  to  inquire  into  the  nature  of  the 
various  kinds  of  changes  and  to  find  some  basis  upon  which  they 
may  be  classified. 

1 


2    -:  :     :  ft  ..*'  GENERAL  CHEMISTRY 

MATTER 

Weight  and  mass.  We  usually  gain  our  impressions  as  to  the  quan- 
tity of  matter  constituting  a  body  by  lifting  it,  or,  in  other  words, 
from  our  muscular  sensation  in  opposing  the  force  of  gravity.  A 
spring  balance  indicates  the  same  quantity  in  a  mechanical  way,  the 
scale  on  such  a  balance  being  constructed  by  marking  the  points  to 
which  a  series  of  arbitrary  units  stretch  the  spring.  What  is  really 
measured  in  this  way  is  the  intensity  of  the  earth's  attraction  for  the 
body,  and  since  this  varies  with  the  distance  of  the  body  from  the 
center  of  the  earth,  it  is  evident  that  the  same  body  will  not  indicate 
the  same  weight  at  places  of  different  elevations.  -The  measure  of 
the  earth's  attraction  for  a  body  is  called  its  weight. 

If  we  employ  an '  ordinary  beam  balance,  counterpoising  the  body 
with  a  set  of  arbitrary  "weights,"  it  will  remain  counterpoised  every- 
where, since  variations  of  gravity  will  affect  both  the  body  and  the 
weights  equally.  The  quantity  of  matter  as  measured  by  being 
balanced  by  arbitrary  units  of  any  kind  is  called  its  mass.  It  is  the 
mass  which  such  a  balance  records,  though  it  is  usually  called  the 
weight.  In  any 'one  locality  the  weight  is  evidently  proportional  to 
the  mass,  and  for  most  purposes  in  chemical  work  the  two  terms  can 
be  employed  interchangeably. 

The  unit  of  mass  which  is  universally  employed  in  scientific  work 
is  the  gram.  This  is  the  one  thousandth  part  of  the  mass  of  a  par- 
ticular piece  of  platinum  preserved  in  the  International  Bureau  of 
Weights  and  Measures  at  Sevres  near  Paris,  and  called  the  standard 
kilogram.  This  mass  was  originally  so  chosen  that  the  mass  of  a 
gram  would  exactly  equal  that  of  one  cubic  centimeter  of  water  at  4° 
centigrade.  Unfortunately  the  effort  to  have  this  so  was  not  entirely 
successful,  and  one  cubic  centimeter  of  water  does  not  have  a  mass 
of  precisely  one  gram,  though  the  difference  is  very  slight. 

Conservation  of  mass.  Since  the  outward  form  and  condition  of 
matter  are  constantly  changing,  it  would  perhaps  seem  reasonable  to 
expect  that  its  mass  is  also  undergoing  variation.  The  mass  of  a 
lump  of  coal  is  certainly  much  greater  than  that  of  the  ashes  which 
result  from  its  burning.  However,  if  we  collect  the  gases  formed  and 
take  into  account  all  the  materials  concerned  in  the  burning,  we  find 
that  the  total  mass  remains  unchanged.  All  of  our  experience  goes 
to  chow  that  this  is  true  in  every  change  in  matter,  and  this  experience 


MATTER  AND  ENERGY  3 

is-_£xpressed  in  the  law  of  conservation  of  mass.  The  law  may  be 
stated  thus :  In  any  change  taking  place  in  a  system  of  materials  the 
mass  remains  constant. 

Density.  The  density  of  a  body  is  the  mass  of  its  unit  volume, 
namely,  of  one  cubic  centimeter.  Since  this  volume  of  water  has  a 
mass  of  almost  exactly  one  gram,  the  figures  which  state  the  den- 
sity of  a  body  at  the  «ame  time  tell  how  many  tunes  as  heavy  as 
water  it  is. 

Gases  are  so  very  light  that  their  densities  must  be  expressed  in 
numbers  which  are  inconveniently  small,  that  of  oxygen  being 
0.001429.  Such  numbers  are  called  the  absolute  densities,  to  distin- 
guish them  from  the  relative  densities  which  are  more  often  employed. 
The  relative  density  is  the  ratio  between  the  mass  of  a  given  volume 
of  one  gas  to  that  of  some  other  gas  chosen  as  a  standard*  For  many 
years  air  was  employed  as  the  standard  gas,  so  that  a  statement  of 
the  relative  density  of  a  gas  was  equivalent  to  a  statement  of  how 
many  times  as  heavy  as  air  it  is.  In  later  years  it  is  coming  to  be  the 
custom  to  compare  gases  with  oxygen  and  to  assign  the  latter  a 
value  of  16  instead  of  unity,  in  order  that  the  density  of  hydrogen, 
the  lightest  of  gases,  may  not  be  less  than  unity.  On  this  basis,  the 
statement  that  the  density  of  nitrogen  is  14.01  means  that,  volume 
for  volume,  the  mass  of  nitrogen  is  to  that  of  oxygen  as  14.01  is 
to  16.  In  any  given  case  the  size  of  the  figures  leaves  no  doubt  as 
to  the  system  used. 

The  properties  of  substances.  By  the  term  properties  is  meant  all 
of  those  marks  or  characteristics  by  which  a  given  substance  is  identi- 
fied. Many  of  these  properties  are  designated  by  words  in  common  use, 
such  as  color,  hardness,  luster,  transparency,  solubility,  melting  point, 
boiling  point,  and  physical  state  (that  is,  whether  solid,  liquid,  or  gas). 
Others  are  not  so  obvious  to  an  untrained  observer,  among  these 
being  density,  refractive  index,  conductivity  for  heat  and  electricity, 
and  crystalline  form. 

All  of  these  properties  are  the .  peculiar  marks  of  the  substance 
itself,  and  are  independent  of  the  presence  or  absence  of  other  sub- 
stances. They  are,  however,  greatly  modified  by  the  physical  con- 
ditions under  which  they  are  observed,  such  as  the  temperature  and  the 
pressure.  Thus  red  oxide  of  mercury  becomes  nearly  black  when 
heated,  but  recovers  its  original  color  when  cooled.  Water  decreases 
in  density  with  rise  of  temperature  as  do  nearly  all  liquids,  so  that 


4  GENERAL  CHEMISTRY 

the  warmer  water  tends  to  rise  to  the  top.  All  gases  may  be  con- 
verted into  liquids  by  pressure  at  low  temperatures,  and  the  boiling 
point,  as  well  as  the  freezing  point,  depends  upon  the  pressure. 

ENERGY 

Since  every  change  in  matter  involves  a  change  in  energy,  we  must 
consider  some  of  the  characteristics  of  energy  before  going  on  to  a 
more  careful  study  of  the  variety  of  changes  in  matter. 

Definition  of  energy.  In  everyday  language  we  sometimes  say  that 
a  certain  man  is  full  of  energy,  meaning  that  he  has  a  great  capacity 
for  wtxrk.  We  also  recognize  the  same  quality  in  inanimate  things. 
We  realize  that  steam,  highly  compressed  in  a  boiler,  possesses  energy, 
for  we  know  that  when  it  is  admitted  to  the  cylinder  of  .the  locomotive 
it  will  push  back  the  piston  head  and  move  the  train.  Energy  is  present- 
in  electric  power  lines,  for  a  motor  connected  with  them  is  caused  to 
rotate,  an  electric  lamp  to  give  light,  a  resistance  heater  to  supply  heat. 
Energy  must  be  expended  in  raising  a  block  of  stone  to  the  top  of  a 
building,  and  when  there  it  too  possesses  energy,  for  by  its  fall  it  will 
produce  effects  which  we  recognize  to  be  due  to  the  expenditure  of 
energy.  We  may  therefore  regard  energy  as  the  ability  to  do  work. 

Varieties  of  energy.  Energy  appears  in  many  forms,  among  the  most 
familiar  of  which  are  heat  energy,  electrical  energy,  and  kinetic  energy, 
or  the  energy  of  moving  bodies.  No  less  important  is  the  potential 
energy  which  a  body  possesses  by  virtue  of  its  position.  This  repre- 
sents the  work  done  upon  a  body  in  raising  it  to  a  height,  and  which 
can  then  be  recovered  in  other  forms  when  the  body  falls. 

Conservation  of  energy.  The  experience  gained  in  a  century  of  ex- 
perimenting has  convinced  scientists  that  it  is  impossible  to  alter  the 
quantity  of  energy  in  a  system  of  bodies,  save  as  we  add  energy  from 
without  or  allow  it  to  escape  from  the  system,  and  this  generalization 
is  known  as  the  law  of  conservation  of  energy.  It  is  not  difficult, 
however,  to  alter  the  distribution  of  energy  between  bodies.  If  a  piece 
of  hot  metal  is  dipped  into  water,  the  metal  is  cooled  and  the  water  is 
heated,  so  that  the  metal  loses  energy  and  the  water  gains  it.  When 
a  swinging  bat  strikes  the  ball,  the  latter  gains  energy  while  the  bat 
loses  it.  There  is  therefore  no  constant  quantity  xrf  energy"  in  a  body 
as  there  is  mass. 

Transformation  of  energy.  Mjetf eover,  energy  can  J>e ,  freely  trans- 
formed from  one  variety  into  another.  .The  heat  energy  of  burning  coal 


MATTER  AND  ENERGY  5 

can  be  changed  into  the  kinetic  energy  of  the  locomotive.  The  kinetic 
energy  of  falling  water  can  be  transformed  into  electrical  energy,  as  in 
the  power  plants  of  Niagara  Falls.  The  electrical  energy  of  the  trolley 
line  is  readily  converted  into  the  kinetic  energy  of  the  moving  car.  In 
all  such  transformations  a  definite  quantity  of  energy  of  one  kind 
always  gives  a  definite  quantity  of  another,  so  that  we  speak  of 
the  mechanical  equivalent  of  heat,  or  the  electrical  equivalent  of 
mechanical  energy. 


FIG.  1 


The  diagram  (Fig.  1)  illustrates  a  few  familiar  transformations  of  energy. 
The  heat  of  the  flame  A  is  converted  into  mechanical  energy  in  the  heat  engine 
E.  The  motion  of  the  engine  is  communicated  to  the  small  dynamo  C,  where  it 
is  converted  into  magnetic  and  electrical  energy.  The  electrical  energy  is  changed 
into  heat  and  light  in  the  incandescent  lamp  D,  and  into  chemical  energy  (see 
following  paragraph)  in  the  electrolysis  of  a  solution  in  the  cell  E. 

Chemical  energy.  Among  the  most  important  kinds  of  energy  is  that 
form  which  is  called  chemical  energy.  In  the  next  chapter  it  will  be 
shown  that  when  substances  burn  in  the  air,  the  change  occurring  is 
due  to  the  action  upon  them  of  an  invisible  gas  called  oxygen.  When 
a  strip  of  magnesium  ribbon  is  burned,  it  is  converted  into  a  white  ash 
and  a  great  deal  of  heat  is  given  off  together  with  an  intense  light. 
Accepting  the  law  of  conservation  of  energy  as  true,  we  argue  that 
this  heat  must  have  come  from  some  other  form  of  energy  present  in 
the  magnesium  and  oxygen.  This  conclusion  is  strengthened  when  we 
find  by  experiment  that  in  order  to  recover  the  magnesium  and  the 
oxygen  from  the  white  ash,  work  must  be  done  equivalent  to  the  heat 
set  free  when  the  -magnesium  burned.  The  latter  can  then  burn  once 
more  with  the  same  evolution  of  he'at  as  in  the  former  experiment. 


.    6  GENERAL  CHEMISTRY 

The  case  is  similar  in  many  respects  to  the  mor'e  familiar  one- pre- 
sented by  falling  bodies.  When  a  body  falls  from  a  height  and  strikes 
the  earth,  its  potential  energy  is  turned  into  mechanical  effects  and 
heat.  To  restore  the  body  to  the  height,  work  must  be  done  upon  it 
equivalent  to  the,  energy  set  free  in  its  fall.  This  work  is  in  some  way 
stored  within  the  body  as  energy,  for  it  may  once  more  fall  with  the 
same  effects  as  before.  We  do  not  know  how  the  energy  is  stored 
within  the  body,  but  we  do  know  that  it  is,  and  we  name  it  potential 
energy.  We  say  that  the  cause  of  the  fall  of  the  body  is  the  attraction 
.of  gravity.  We  can  measure  the  force  of  gravity,  and  we  know  the 
law^_^e^iilatirigthe  fall  of  bodies,  butwe  do  not  understand  the 
natma-o£  .^ravitynor  how  ItTiu**0  ^ 

In  a  similar  way  we  give  the  name  Chemical  energy  to  that  form 
of  energy  which  enables  substances  to  combine  with  each  other,  and 
we  say  that  they  combine  because  of  chemical  attraction  or  affinity; 
but  we  do  not  know  anything  more  about  the  nature  oi1  chemical 
energy  than  we  do  about  potential  energy,  nor  do  we  know  anything 
more  about  chemical  affinity  than  we  know  about  gravitational  attrac- 
Jdon.  These  terms  are  merely  names  for  things  which  we  know  to  exist, 
but  whose  nature  is  as  yet  unknown  to  us. 

The  measurement  of  energy.  Since  changes  in  energy,  both  from 
one  body  to  another  and  from  one  form  into  another,  are  invariably 
involved  in  chemical  action,  it  is  a  matter  of  great  practical  impor- 
tance to  devise  units  for  the  measurement  of  energy,  and  methods  for 
making  the  measurement.  In  general  each  kind  of  energy  must  have 
its  own  units,  just  as  with  matter  we  have  centimeters  for  lengths, 
liters  for  volumes,  grams  for  weights.  Tn__^nnip  nf  ifg  fojrps  pTipygy  1g 
very  difficult  to  measure  directly,  and  neither  units  nor  methods  for 
the  direct  measurement  of  chemical  eHergyliave  as  yet  been  devised. 
In  such  cases  it  is  necessary  to  transform  the  energy  into  a  form  more 
pnn-\mniftj^^fQr^measurement.^  In  the  case  of  chemical  energy  it  is 
changed  into  heat  or  electrical  energy  for  this  purpose. 

Measurement  of  heat.  The  thermometer  measures  the  intensity. -of 
^&oak? — how^hot  a  body  is^^  not  the  quantity  of  its  heat  energy.  A 
very  fine  platinum  wire  may  be  readily  heated  to  1600°  or  above,  but 
the  actual  heat  which  it  will  give  out  on  cooling  is  very  small.  JTftat. 
energy  is  measured  by  observing  to  what  extent  it  will  change  the 
temperature  of  a  given  mass  of  some  standard  substance.  Water  has 
been  chosen  as  the  standard,  and  the  unit  of  heat,  called  the  calorie 


MATTER  AND  ENERGY 


(designated  by  the  abbreviation  raZ.),  is  denned  as  the  quantity  re- 
quired to  change  the  temperature  of  a  gram  of  water  one  degree  on 
the  centigrade  scale.  Sometimes  a  larger  unit  is  desirable,  and  this  is 
taken  as  1000  times  the  unit  just  denned.  This  is  called  the  large 
calorie  and  is  designated  by  the  abbreviation  Cat. 

The  actual  measurement  of  the  quantity  of  chemical  energy  transformed  into 
heat  in  a  chemical'  action  is  accomplished  by  the  use  of  an  apparatus  called 
Ihe  calorimeter,  represented  in  Fig.  2.  The  action  is  arranged  to  take  place  in 
solution  in  a  measured  volume  of  water  contained  in 
a  thin-walled  metal  vessel  A.  This  is  placed  within  a 
double-walled  vessel  B,  which  contains  water  at  the 
temperature  of  the  room.  The  thermometer  C  indi- 
cates when  the  water  has  reached  this  temperature. 
This  is  to  prevent  the  influence  of  heat  from  without, 
and  as  an  added  precaution  the  vessel  is  covered  with 
a  thick  layer  of  nonconducting  felt.  The  heat  evolved 
by  the  reaction  raises  the  temperature  of  the  solution, 
the  rise  being  indicated  by  the  thermometers  D,D. 
During  the  reaction  the  solution  is  stirred  by  the 
stirrer  E.  If  the  weight  of  the  water  is  2570  g.  and 
the  rise  in  temperature  is  1.5°,  the  heat  evolved  is 
2570  x  1.5  =  3855  cal. 


FIG.  2 


VARIETIES  OF  MATTER 


The  variety  of  forms  which  matter  assumes  in  all  the  wonderful 
transformations  of  nature  is  almost  infinite,  and  these  may  be  classified 
in  a  great  many  ways,  according  to  the  purpose  in  view.  The  inter- 
est of  the  chemist  centers  chiefly  in  the  composition  of  substances  and 
in  their  chemical  energy,  together  with  the  changes  which  take  place 
in  both  of  these.  From  this  standpoint  he  finds  it  convenient  to  place 
them  in  three  groups;  namely,  compounds,  elements,  and  mixtures. 
The  distinction  can  be  best  explained  by  some  experiments. 

Illustrative  experiments.  The  chief  characteristics  of  the  substances 
iron  and  suJfur  are  familiar  to  almost  every  one.  Iron  filings  form  a 
heavy  black  powder,  strongly  attracted  by  a  magnet.  When  treated 
with  the  liquid  known  as  hydrochloric  acid,  the  iron  passes  into  solu- 
tion and  a  colorless  gas  called  hydrogen  is  evolved,  considerable  heat 
being  liberated  in  the  process.  Sulfur  may  be  obtained  as  a  light  yellow 
powder  not  attracted  by  a  magnet  nor  dissolved  by  hydrochloric 
acid.  It  is  readily  soluble  in  the  liquid  known  as  carbon  disulfide, 
however,  which  is  not  true  of  iron,  and  when  the  solution  is  allowed 
to  evaporate,  the  sulfur  is  deposited  in  the  form  of  yellow  crystals. 


GENERAL  CHEMISTRY 


When  these  two  materials  are  thoroughly  ground  together,  a  green- 
ish-black powder  is  obtained  which  is  quite  different  from  either  of 
them ;  but  when  we  apply  the  tests  which  we  have  found  to  char- 
acterize iron  and  sulfur,  it  is  found  that  in  many  respects  it  acts 
like  these  two  taken  separately.  Hydrochloric  acid  still  dissolves  the 
iron  and  evolves  hydrogen  with  the~same  heat  as  before,  leaving  the 
sulfur  unchanged.  A  magnet  rubbed  through  it  withdraws  the  iron 
and  leaves  the  sulfur.  Carbon  disulfide  dissolves  the  sulfur  but  not 
the  iron.  The  sulfur  and  the  iron  each  act  just  as  they  did  before 
they  were  ground  together,  and  with  the  same  energy. 

If  now  a  portion  of  the  powder  is  placed  in  a  test  tube  and  heated, 
as  shown  in  Fig.  3,  it  soon  begins  to  glow  at  the  point  of  greatest 

heat,  and  even  if  the  flame  is  with- 
drawn, the  glow  continues  to  spread 
throughout  the  entire  contents  of 
the  test  tube,  and  a  great  deal  of 
heat  is  set  free  at  the  same  time. 
When  the  product  is  examined,  it 
is  found  that  many  of  the  character- 
istics of  the  iron  and  sulfur  have 
been  modified.  Carbon  disulfide  no 
longer  dissolves  sulfur  and  leaves 
iron ;  a  magnet  has  no  effect  upon 

the  material ;  hydrochloric  acid  dissolves  the  entire  product  and 
evolves  a  gas  of  disagreeable  odor  quite  different  from  hydrogen; 
and  the  heat  liberated  is  quite  different  in  quantity  from  that  in 
the  former  case.  Many  of  the  properties  of  this  product  differ  from 
those  of  the  constituents,  among  these  being  density,  color,  hardness, 
solubility,  and  melting  point.  The  material  is  called  iron  sulfide. 

When,  a  small  quantity  of  sugar  is  heated  in  a  test  tube,  it  melts, 
turns  brown  in  color,  gives  off  vapors,  and  finally  dries  up  to  a  solid 
black  residue  which  may  be  identified  as  carbon.  By  collecting  and 
examining  the  vapors  it  is  possible  to  show  that  they  are  largely  water. 
To  make  this  transformation  complete  it  is  necessary  to  apply  heat 
throughout  the  entire  process.  In  this  experiment  one  substance,  merely 
by  the  application  of  heat,  has  given  rise  to  at  least  two  others  of  very 
different  properties,  and  the  change  is  described  as  a  decomposition. 

Chemical  reactions.  The  two  examples  which  have  just  been  de- 
scribed are  very  different  in  many  ways,  but  they  have  several 


FIG.  3 


MATTER  AXD   ENERGY  9 

characteristics  in  common.  Most  of  the  properties  of  the  materials  con- 
cerned undergo  a  very  considerable  change,  so  that  it  is  easy  to  recog- 
nize the  fact  that  the  products  formed  are  different  from  the  original 
materials.  A  more  important  characteristic  is  that  the  chemical  energy 
of  the  materials  has  been  changed.  The  action  of  iron  with  sulfur, 
when  once  started,  is  attended  by  the  evolution  of  a  great  deal  of 
heat  which  is  at  the  expense  of  chemical  energy.  That  the  iron  sulfide 
contains  less  chemical  energy  is  shown  by  the  fact  that  work  must  be 
done  upon  it  to  regain  the  iron  and  sulfur,  and  that  it  is  not  able  to 
act  upon  other  materials  as  do  iron  and  sulfur  individually.  To  effect 
the  decomposition  of  sugar  it  is  necessary  to  supply  heat  during  the 
entire  period  of  decomposition,  and  this  heat  must  be  converted  into 
some  other  form  of  energy.  A  part  of  it  is  converted  into  chemical 
energy,  so  that  the  carbon  and  the  water  taken  together  represent 
more  energy  of  this  kind  than  does  the  original  sugar.  Any  trans- 
formation in  matter,  whether  union  or  decomposition,  which  involves 
a  change  in  the  chemical  energy  of  the  substances  concerned  is  called 
a  chemical  reaction. 

Definition  of  chemical  compounds.  When  it  can  be  shown  that  a 
substance  is  composed  of  at  least  two  different  materials,  and  that  its 
"chemcal  energy  is  different  from"  that  of  thj  constituents  iakeiTsep- 
aratelyTitTis  called  a  chemical  compound.  Thus  we  judge  iron  sulfide 
to  be  a  chemical  compound,  tor  it  contains  iron  ana  sulfur,  and  yet 
differs  from  them  in  chemical  energy.  It  will  be  seen  that  in  some 
cases  this  difference  is  so  slight  that  it  may  be  impossible  to  tell 
whether  there  is  a  real  difference  or  not,  especially  since  it  is  difficult 
to  measure  chemical  energy  with  accuracy  even  indirectly.  It  is  there- 
fore not  always  possible  to  decide  whether  a  given  material  is  a 
chemical  compound  or  not.  We  are  assisted  in  our  decision  by  the 
fact,  to  be  proved  in  a  later  chapter,  that  the  percentage  composition 
f  a  true  compound  is  always  the  same..  If  the  material  can  be  ob- 
tained in  pure  form  and  analyzed,  it  is  usually  possible  to  decide  the 
question  in  this  way. 

Other  energy  changes  involved  in  chemical  action.  Since  in  a  chemi- 
cal reaction  there  is  always  a  change  in  chemical  energy,  it  is  evident 
that  there  must  be  other  energy  changes  as  well.  Usually  the  loss  or 
gain  in  chemical  energy  is  indicated  by  a  corresponding  evolution  or 
absorption  of  heat,  as  in  the  experiments  with  iron  sulfide  and  sugar. 
Under  the  proper  conditions  the  chemical  energy  may  be  converted 


10  GENERAL  CHEMISTRY 

into  electrical  energy.  Thus  if  a  plate  of  zinc  is  connected  by  a  wire 
with  one  of  carbon,  and  the  two^are  dipped  into  hydrochloric  acid,  the 
zinc  is  acted  upon  by  the  acid,  and  the  chemical  energy  set  free  ap- 
pears as  an  electric  current  in  the  wire.  In  other  cases  it  is  partially 
converted  into  light,  as  when  magnesium  wire  burns,  or  into  kinetic 
energy,  as  when  the  chemical  action  of  gunpowder  imparts  motion  to 
the  bullet. 

It  is  by  no  means  true,  however,  that  a  change  in  energy  may  be 
considered  as  evidence  of  chemical  action.  As  we  have  seen,  chemical 
energy  is  op1y  ™IP.  ryf  jnany  forms  which  may  be  transformed  into  the 
familiar  forms  of  heat  and  electrical  energy.  .Therefore  an  energy 
change  in  a  given  case  may  be  due  to  one  or  more  of  many  causes 
other  than  chemical  action.  The  energy  which  we  observe  on  turning 
the  key  to  an  electric  light  is  due  to  the  mechanical  energy  of  the 
engine,  which  may  be  driven  by  water  power  and  involve  no  chemical 
action  at  all.  The  heat  generated  when  a  gas  is  compressed  represents 
only  the  mechanical  work  done  upon  the  gas. 

Conditions  affecting  chemical  action.  There  are  many  conditions 
which  may  either  promote  or  hinder  chemical  action'.  An  increase  in 
temperature  is  usually  favorable  to  chemical  action,  as  was  seen  in 
the  case  of  iron  and  sulfur.  It  frequently  promotes  decomposition,  as 
in  the  case  of  sugar.  Other  forms  of  energy,  such  as  light,  mechanical 
pressure,  shock,  and  electrical  energy,  may  also  facilitate  either  chemi- 
cal union  or  decomposition,  at  times  overcoming  obstacles  which  pre- 
vent union,  in  other  cases  overpowering  the  chemical  attraction  which 
holds  a  compound  together. 

Chemical  conduct  of  substances.  We  have  seen  that  substances  are 
characterized  by  their  properties.  They  are  also  characterized  by  the 
way  in  which  they  act  toward  other  substances,  as  well  as  by  the 
decompositions  which  they  undergo  under  various  conditions.  Thus 
a  substance  may  burn  with  a  flame  when  heated  in  the  air,  or  it  may 
combine  with  another  substance  with  incandescence,  as  is  the  case  of 
iron  heated  with  sulfur.  It  may  decompose  when  heated,  as  is  true 
of  sugar,  or  when  subjected  to  the  action  of  the  electric  current,  as 
water  does.  It  may  have  no  action  at  all  with  certain  substances 
under  ordinary  circumstances,  while  with  some  of  them  it  may  combine 
at  high  temperatures  and  with  others  remain  indifferent.  All  such 
peculiarities  are  collectively  called  the  chemical  conduct  of  a  sub- 
stance. It  will  be  seen  that  such  conduct  usually  depends  upon  the 


MATTER  AND  ENERGY 


11 


presence  of  other  materials,  as  well  #s  upon  the  physi^aj  conditions 
which  prevail,  such  as  temperature  and  pressure. 

Elements.  The  experiments  outlined  suggest  that  we  heat,  iron, 
sulfur,  water,  and  carbon  and^ee  whether  any  new  substance  can  be 
obtained  through  their  decomposition.  Experience  has  shown  that  in 
many  cases  decompositions  may  be  brought  about  by  the  electrical 
current  or  by  the  action  of  substances  possessed  of  great  chemical 
energy,  and  we  may  also  employ  these 
methods.  In  such  ways  chejnisto  have  OUG* 
ceeded  in  decomposing  water  into  two 
invgibte^ases,  Oxygen  and  hydrogen.  scT 
that  it  must  be  regarded  as  a  compound. 

The  decomposition  of  water.  The  decomposition 
of  water  may  readily  be  observed  by  the  aid  of  an 
apparatus  such  as  that  represented  in  Fig.  4.  Two 
test  tubes  (A  and  B)  are  filled  with  water  and 
inverted  in  a  vessel  half  filled  with  water  to  which 
a  little  sulfuric  acid  has  been  added.  A  piece  of 
platinum  foil  (C  and  Z>)  attached  to  a  wire  is  then 
brought  under  the  end  of  each  tube.  When  these 
wires  are  connected  with  a  source  of  current  sup- 
plying from  6  to  10  volts,  bubbles  of  gas  will  be  FIG.  4 

seen  to  form  in  each  tube.    These  gases  may  be 

shown  to  have  different  properties  ;  they  are  hydrogen  and  oxygen.  The  reason 
for  adding  the  sulfuric  acid  cannot  be  explained  at  this  time,  but  will  be 
discussed  later  on. 


%J    Vg-^- 


The  other  three  anbatfl/nnp-a  .  p.a.rhnn1  i 


snlfrir.  havp.  never  been 


decomposed,  notwithstanding  the  many  efforts  directed  to  this  end. 
Substances  like  these  three,  which  have  never  been  decomposed  into 
two  or  more  different  materials,  are  called  elements,  or  elementary  sub- 
stances. {t  should  be  carefully  noted,  however,  that  this  definition 
does  not  suggest  anything  as  to  the  real  nature  of  an  element.  Neither 
does  it  preclude  the  possibility  that  one  element  may  be  transformed 
into  another  under  some  condition  which  changes  its  chemical  energy. 
Indeed,  there  is  reason  for  believing  that  all  the  elements  have  had 
a  common  origin,  and  in  the  discussion  of  radium  it  will  be  shown 
that  this  element  at  least  is  slowly  changing  into  others. 

Number  of  the  elements.  While  many  thousands  of  compounds 
have  been  described,  the  number  of  the  elements  at  present  known 
is  comparatively  small,  a  complete  list  being  given  in  the  table  on 
the  following  page.  There  is  good  reason  for  thinking  that  no  very 


12 


GENERAL   CHEMISTRY 


TABLE  OF  THE   ELEMENTS 


ELEMENT 

SYMBOL 

ATOMIC 
WEIGHT 

ELEMENT 

SYMBOL 

ATOMIC 
WEIGHT 

Aluminium   .     .     . 

Al 

27.1 

Neodymium    . 

Nd 

144.3 

Antimony      .     .     . 

Sb 

120.2 

Neon      .... 

Ne 

20.2 

Araron 

A 

39.88 

Niton 

Nt 

222.4 

Arsenic     .... 

As 

74.96 

Nickel    .... 

Ni 

58.68 

Barium     .... 

Ba 

137.37 

Nitrogen     .     .     . 

N 

14.01 

Bismuth   .... 

Bi 

208.0 

Osmium      .     .     . 

Os 

190.9 

B 

11.0 

Oxvsren 

o 

16.00 

Bromine   .... 

Br 

79.92 

V^-A-jr^t/JJ.    ... 

Palladium  .     .     . 

Pd 

106.7 

Cadmium 

Cd 

112.40 

Phosphorus 

P 

31.04 

Caesium     .... 

Cs 

132.81 

Platinum    . 

Pt 

195.2 

Calcium    .... 

Ca 

40.07 

Potassium  .     .     . 

K 

39.10 

Carbon      .... 

C 

12.00 

Praseodymium     . 

Pr 

140.6 

Cerium      .... 

Ce 

140.25 

Radium      .     .'    . 

Ra 

226.4 

Chlorine   .... 

Cl 

35.46 

Rhodium    . 

Rh 

102.9 

Chromium 

Cr 

52.0 

Rubidium  . 

Rb 

85.45 

Cobalt       .... 

Co 

58*.97 

Ruthenium     ..     . 

Ru 

101.7 

Columbium  . 

Cb 

93.5 

Samarium  . 

Sa 

150.4 

Copper      .     .     .     . 

Cu 

63.57 

Scandium   . 

Sc 

44.1 

Dysprosium  .     .     . 

Dy 

162.5 

Selenium    . 

Se 

79.2 

Erbium     .... 

Er 

167.7 

Silicon   .... 

Si 

28.3 

Europium 

Eu 

152.0 

Silver     .... 

Ag 

107.88 

Fluorine   .... 

F 

19.0 

Sodium  .... 

Na 

23.00 

Gadolinium  . 

Gd 

157.3 

Strontium  . 

Sr 

87.63 

Gallium    .... 

Ga 

69.9 

Sulfur    .... 

S 

32.07 

Germanium  . 

Ge 

72.5 

Tantalum  .     .     . 

Ta 

181.5 

Glucinum 

Gl 

9.1 

Tellurium  .     .     . 

Te 

127.5 

Gold     .     . 

Au 

107  9 

Tprbinm 

Th 

1  KQ    O 

Helium     .... 

He 

-L  t7  t  •  £ 

3.99 

1   *    I  I  M  U  1  1  1          ... 

Thallium    .     .     . 

A.  D 

Tl 

204.0 

Hydrogen      .     .     . 

H 

i.oofc 

Thorium     . 

Th 

232.4 

Indium     .... 

In 

114.8 

Thulium     .     .     . 

Tm 

168.5 

Iodine       .... 

I 

126.92 

Tin   ... 

Sn 

119.0 

Iridium     .... 

Ir 

193.1 

Titanium    .     .     . 

Ti 

48.1 

Iron      

Fe 

55.84 

Tnnoxtpn 

1  &J_  n 

-i.  M  1  1^  >l  '    II       ... 

.Lo'r.U 

Krypton    .... 

Kr 

82.9 

Uranium     .     .     . 

U 

238.5 

Lanthanum  . 

La 

139.0 

Vanadium  . 

V 

51.0 

Lead    

Pb 

207.10 

y 

•^r 

i  ^o  9 

Lithium    .... 

Li 

6.94 

Ytterbium            ^ 

Lutecium 

Lu 

174.0 

(Neoytterbium)  j 

Yb 

172.0 

Magnesium   . 

Mg 

24.32 

Yttrium      .     .     . 

Yt 

89.0 

Manganese    . 

Mn 

54.93 

Zinc 

Zn 

65.37 

Mercury   .... 

Hg 

200.6 

Zirconium  .     .     . 

Zr 

90.6 

Molybdenum      .     . 

Mo 

96.0 

MATTER  AND  ENERGY 


13 


considerable  number  will  ever  be  added  to  this  list,  and  while  it  is  pos- 
sible  that  a  few_of  t.h™?A  nnw  ligfced  will  turn  out  to  be  compounds  or 
mixtures^  it  is  probable  that  mnsf,  of  them  are  really  the  elementary 
materials  out  of  which  all  others  are  composed.  Following  the  name  of_ 
^agh  element  ill  the  table  is  an  abbreviation  called  a  symbol^  by  which 
the  element  is  designated  among  chemists.  The  symbol  is  usually  the 
initial  letter  of "  the" name  of  the  element,  together  with  some  other 
characteristic  letter.  In  the  case j3f_sjam&-qf  the  elements  the  symbol 
is  the  abbreviation  of  the  old  Latin  name,  as  is  true  of  iron 
gold(awrMw),  and  mercury  (hydrargyrum).  The  significance  of  the 
column  of  numbers  will  be  made  clear  a  little  later. 


Distribution  of  the  elements.  So  far  as  we  can  judge,  these  elements 
are  of  very  unequal  occurrence  in  nature.  It  must  be  remembered, 
however,  that  our  knowledge  of  the  earth's  composition  is  confined 
to  what  is  comparatively  a  very  thin  surface  shell,  not  exceeding  a 
few  miles  in  thickness.  The  table  on  page  14,  prepared  by  F.  W. 
Clarke  and  based  on  the  analysis  of  representative  rocks  and  min- 
erals, gives  an  estimate  of  the  composition  of  this  shell.  It  will  be 
seen  that  20  of  the  elements  are  estimated  to  constitute_99.5  per  cent 
of  the  shell,  the  other  60  together  making  up  the  remaining  0.5  per^ 
cent.  Some  of  the  elements  are  of  such  rare  occurrenceTJiat  only  a 
few  grains  have  ever  been  isolated^ 

Elements  essential  to  life.  A  careful  examination  of  the  materials 
present  in  living  organisms  shows  that  only  a  very  few  are  of  vital 
importance  to  us.  The  following  table,  compiled  by  Sherman,  indi- 
cates the  average  composition  of  the  human  body.  It  is  possible  that 
other  elements  have  an  importance  which  we  do  not  realize,  but,  so  far 
as  we  can  judge,  these  are  the  only  ones  upon  which  living  organisms 


AVERAGE  COMPOSITION  OF  THE  HUMAN  BODY 


Oxygen 
Carbon  . 
Hydrogen 
Nitrogen 
Calcium 


65.00% 

18.00% 

10.00% 

3.00% 

2.00% 


Phosphorus 
Potassium  . 
Sulfur  .  . 
Sodium .  . 
Chlorine  , 


1.00% 
0.35% 
0.25% 
0.15% 
0.15% 


Magnesium 
Iron .  .  . 
Iodine  .  . 
Fluorine  . 
Silicon  . 


0.05% 

0.004% 

traces 

traces 

traces 


Mixtures.  It  is  quite  possible  to  prepare,  from  either  elements  or 
compounds,  or  both,  a  body  which  is  not  itself  a  compound  but  is 
merely  a  mixture.  Ordinary  concrete  is  such  a  material,  for  in  a 


14 


GENERAL  CHEMISTRY 


CLARKE'S  TABLE 


SOLID  SHELL 
93  PER  CENT 

OCEAN  7  PER 
CENT 

AVERAGE 
INCLUDING 
ATMOSPHERE 

47.07 

85.79 

49.78 

Silicon 

28.06 

26.08 

Aluminium 

7.90 

7.34 

4.43 

4.11 

Calcium 

3.44 

0.05 

3.19 

Miasrnesi  uni 

240 

0.14 

2.24 

Sodium                                   • 

243 

1.14 

233 

2.45 

0.04 

2.28 

Hydrogen     

0.22 

10.07 

0.95 

0.40 

0.37 

Carbon               ... 

0.20 

0.002 

0.19 

0.07 

2.07 

0.21 

Bromine                                 . 

0.008 

0.11 

0.11 

Sulfur      

0.11 

0.09 

0.11 

0.09 

0.09 

Manganese  ....... 

0.07 

0.07 

Strontium          ... 

0.03 

0.03 

Nitrogen 

0.02 

Fluorine 

0.02 

0.02 

Other  elements      

0.50 

0.48 

100.00 

100.00 

100.00 

broken  piece  it  is  easy  to  identify  the  crushed  stone,  the  sand,  and 
the  cement  which  compose  it.  Granite  is  a  sort  of  natural  concrete,  in 
which  two  very  different-looking  crystalline  materials,  mica  and  feld- 
spar, are  bound  together  by  a  glassy  substance  called  silica.  A  crys- 
tal of  mica  broken  out  from  granite  has  exactly  the  same  chemical 
energy  as  does  a  pure  crystal  obtained  from  any  other  source.  Iron 
and  sulfur  when  rubbed  together  form  a  material  more  closely  resem- 
bling a  compound,  in  that  it  is  apparently  of  even  quality  through- 
out, or  is  homogeneous.  An  examination  under  the  microscope  shows 
that  this  is  not  really  so,  for  the  particles  of  iron  and  sulfur  can  still 
be  seen  lying  side  by  side  unchanged. 

In  many  cases  the  two  types,  mixtures  and  compounds,  approach 
so  nearly  to  each  other  that  it  is  impossible  to  distinguish  between 
them.  For  example,  alcohol  and  water  mix  in  all  proportions  to 
form  a  perfectly  homogeneous  liquid ;  copper  ami-sine  when  melted 


MATTER  AND  ENERGY  15 

-together  ^nr1  hra.<?sT  which  iii  properties  is  quite  different  from  either, 
yet  has  no  fixed  percentage  composition.  A  perfectly  definite  com- 
pound, such  as  water,  may  be  regarded  as  standing  at  one  end  of  a 
series,  with  an  undoubted  mixture  like  granite  at  the  other.  There  is 
every  gradation  between  these  two  extremes,  and  in  the  middle  of  the 
series  the  one  type  changes  by  almost  imperceptible  stages  into  the 
other.  In  the  majority  of  cases  the  distinction  is  easily  made,  and 
affords  a  useful  basis  of  classification. 

Plan  of  study.  We  shall  now  take  up  the  study  of  two  abundant 
and  typical  elements,  namely,  oxygen  and  hydrogen,  and  following 
this  a  consideration  of  the  compounds  which  they  form  with  each 
other.  This  will  serve  to  emphasize  the  characteristics  of  chemical 
action,  and  provide  a  basis  for  the  consideration  of  other  elements 
and  their  compounds. 


CHAPTER  II 

OXYGEN 

History.  JosephPriestley,  an  English  clergyman  and  investigator, 

is  usually  regarded  as  the  discoverer  of  oxygen.  In  1774,  in  the 
course  of  some  experiments  with  gases,  or  "  airs,"  as  he  called  them, 
it  occurred  to  him  to  try  the  effect  of  heat  upon  certain  solids,  to  deter- 
mine whether  any  gas  is  liberated  in  the  process,  and,  if  so,  to  collect 
the  gas  and  study  its  properties.  He  introduced  into  a  glass  tube 
a  small  amount  of  the  compound  now  known  as  mercuric  oxide,  and 
heated  it  by  means  of  a  large  lens  used  as  a  burning  glass.  He  found 
that  under  these  conditions  a  colorless  gas  was  set  free,  which  aroused 

%his  interest,  because  "  a  candle  burned  in  this  air  with  a  remarkably 
vigorous  flame."  It  is  now  known  that  the  Swedish  chemist  Scheele 

Jaad  obtained  this  same  gas  a  year  earlier  by  heating  niter  as  well 
as  the  mineral  pyrolusite,  but  Priestley  was  the  first  to  publish  the 
results  of  his  experiments,  and  it  was  through  his  description  of 
the  gasthat  it  becanis^known  to  scienttgtis.^x 

At  the  time  of  Priestley's  discovery  the  renowned  French  chemist 
Lavoisier,  whose  life  was  later  sacrificed  in  the  French  Revolution, 
was  engaged  in  a  study  of  the  nature  of  combustion,  and  he  became 
at  once  greatly  interested  in  this  newly  discovered  gas.  He  found 
that  a  number  of  elements  such  as  phosphorus  and  sulfur  unite  with 
it  to  form  compounds  which  .were  at  that  time  regarded  as  acids. 
Believing  that  the  characteristic  properties  of  acids  were  due  to  the 
presence  in  them  of  this  substance,  he  proposed  for  it  the  name 
oxygen,  a  word  derived  from  the  Greek  and  meaning^  acid  former." 
We  now  know  thattEisTiaTne  iT3~"riot  appropriate,  since  many  acids 
do  not  contain  oxygen. 

Occurrence.  Oxygen  occurs  in  nature  both  as  free  oxygen  and  as 
a  constituent  of  many  compounds.  In  100  volumes  of  dry  air  there 
are  approximately  21  volumes  of  the  free  element.  In  the  combined 
state  it  constitutes  88.81  per  cent  by  weight  of  water  and  nearly  one 
half  by  weight  of  the  common  minerals  such  as  limestone,  sandstone, 
granite,  and  clay,  which  together  make  up  the  earth's  crust.  It  is  also 

16 


OXYGEN 


IT 


an  essential  constituent  of  the  compounds  present  in  living  organisms. 
For  example,  nearly  two  thirds  of  the  human  body  is  oxygen  (see 
table,  p.  13).  The  total  weight  of  oxygen  in  the  land,  the  water,  the 
atmosphere,  and*"1'"  lJYin(°f  ™>ffn-"iaTP*!l  TTW  ^  rftfl-ajdad  as  very  roughly 
equal  to  the  united  weights  of  all  the  other  elements. 

Preparation.  Since  oxygen'Ts  HO  Hfoiindaiit  and  is  present  in  such 
a  large  variety  of  compounds,  it  is  easy  to  understand  why  many 
different  methods  may  be  used  in  obtaining  it  in  pure  condition.  The 
most  important  of  these  methods  are  the  following  : 

1.  By  heating  certain  compounds  of  oxygen.  Many  compounds  con- 
taining oxygen  give  off  at  least  a  portion  of  it  when  heated.  For 

potassium  chlorate  —  compounds  which 


contain  respectively  7.4  percent  and  39.2  per  cent  of  oxygen  —  give 
off  jill^  frhair  me  yg-np_jHrhfin  freat^d  to  a  moderately  hightemperature. 
Other  cormoound&such  as  manganese  dioxide  and  barium  peroxide, 
give.  JipjjnlY  a^  definite  fraction  of  their  oxygen. 

2.  By  the  action  of  certain  compounds  upon  others  rich  in  oxygen.  When 
sulfuric  acid  acts  upon  potassium  dichromate  or  potassium  perman- 
ganate,  both  of   which   compounds  are  rich  in 

oxygen,  a  definite  percentage  of  the  oxygen 
present  is  liberated.  In  a  sjrnjlaj^-wa^-when 
water  acts  upon  sodium  peroxide  under  suitable 
conditions,  oxygen  "is  set  free. 

3.  By  the  decomposition  of  water  by  means  of 
an  electric  current.    It  will  be  shown  in  a  later 
chapter  that  water  is  a  compound  consisting  of 
88.81  per  cent  oxygen  and  11.19  per  cent  hydro- 
gen.  When  pure  it  is  practically  a  nonconductor 
of  electricity,  but  if  a  little  sulfuric  acid  or  potas- 
sium hydroxide  is  added  to  it,  the  resulting  solu- 
tion very  readily  conducts  the  current.    In  this 
process    a  series  of  changes   take   place    which 
result  in  the  decomposition  of  the  water  into  its 
constituents. 


.  5 


Fig.  5  represents  a  convenient  form  of  apparatus  for 
effecting  the  decomposition  of  water  in  this  manner. 
Two  platinum  wires  (A  and  B),  each  with  a  small  piece 
of  platinum  foil  attached  to  one  end,  are  fused  through 

the  tubes  C  and  D,  as  shown  in  the  figure.    The  stopcocks  at  the  tops  of  these 
tubes  are  opened,  and  water,  to  which  has  been  added  about  one  tenth  of  its 


18  GENERAL  CHEMISTRY 

volume  of  sulf  uric  acid,  is  poured  into  the  tube  E  until  the  side  tubes  C  and  D 
are  completely  filled.  The  stopcocks  are  then  closed.  The  platinum  wires  are 
now  connected  with  a  battery  capable  of  supplying  about  6  volts  (3  storage 
or  6  dichromate  cells  joined  in  series).  The  current  flows  through  the  acidu- 
lated water  from  one  piece  of  platinum  foil  (electrode)  to  the  other,  and  brings 
about  the  decomposition  of  the  water  into  its  constituents.  The  oxygen  rises  in 
bubbles  from  the  positive  electrode  and  collects  in  the  upper  part  of  tube  C, 
while  the  hydrogen  rises  from  the  negative  electrode  and  collects  in  tube  D. 

4.  By  separation  from  air.  Since  air  contains  such  a  large  percentage 
of  free  oxygen,  one  would  naturally  expect  methods  to  be  devised  for 
obtaining  it  from  this  source.  The  problem  is  not  as  simple  as  it  may 
seem,  for  there  are  other  gases  in  the  ah-,  and  the  separation  of  a  gas 
in  a  pure  condition  from  a  mixture  of  gases  is  always  difficult.  In 
this  case  it  may  be  accomplished  by  either  of  two  methods : 

(a)  Chemical  method.  A  few  substances  are  known  which,  when 
heated  in  the  air  to  a  temperature  varying  with  the  nature  of  the 
substance,   combine  with  oxygen  present  in  the  air,  but  give  it  up 
once  more  at  a  higher  temperature.    Barium  oxide,  a  compound  con- 
taining 10.43  per  cent  of  oxygen,  is  a  good  example  of  this  kind. 
When  heated  in  the  air  to  a  temperature  of  about  500°,  it  com- 
bines with  oxygen  and  is  thereby  changed  into  barium  peroxide,  a 
compound  which  contains  18.89  per  cent  of  oxygen.    When  this  is 
further  heated  to  about  1000°,  the  additional  oxygen  is  set  free  once 
more  and  may  be  collected,  the  barium  peroxide  being  at  the  same 
time  changed  into  the  oxide  again.    These  transformations  may  be 
represented  thus: 

barium  oxide  +  oxygen  <    >  barium  peroxide 

The  arrows  indicate  that  the  reaction  may  proceed  in  either  direc- 
tion, according  to  the  temperature. 

(b)  Mechanical  method.  By  subjecting  air  to  the  combined  effects 
of  pressure  and  very  low  temperature  it  is  possible  to  obtain  it  in  the 
form  of  a  liquid  which  is  essentially  a  mixture  of  oxygen  and  nitrogen. 
When  this  liquid  is  allowed  to  stand  under  ordinary  pressure,  it  boils 
rapidly,  the  temperature  falling  to  a  very  low  point.    Since  nitrogen 
has  the  lower  boiling  point  (—  195.7°),  it  tends  to  boil  away  first,  and 
is  gradually  followed  by  the  oxygen  (boiling  point  —  182.9°),  which 
may  be  collected  separately.    The  oxygen  prepared  by  this  method 
generally  contains  a  small  percentage  of  nitrogen. 

Practical  methods  of  preparation.  With  these  general  methods  of 
preparation  before  us,  we  may  make  a  selection  of  those  best  suited 


OXYGEN 


19 


FIG.  6 


for  the  actual  preparation  of  the  gas.  For  the  purpose  of  laboratory 
experiments,  in  which  relatively  small  quantities  are  desired,  the 
choice  will  naturally  be  guided  by  convenience  and  simplicity  of  ap- 
paratus, while  in  the  preparation  on  a  commercial  scale  economy  will 
determine  the  method. 

Laboratory  method.  The  method  usually  chosen  for  preparing  oxy- 
gen in  the  laboratory  consists  in  heating  potassium  chlorate,  which 
is  a  white,  solid  compound  of  potassium,  chlo- 
rine, and  oxygen.  The  evolution  of  the  gas 
becomes  marked  at  about  400°,  and  if  the 
heating  is  continued  long  enough,  all  the  oxy- 
gen present  in  the  chlorate  is  liberated.  It  is  a 

remarkable  fact  that 
the  rate  at  which  the 
oxygen  is  evolved 
at  any  given  tem- 
perature is  greatly 
increased  by  the 
presence  of  small 
quantities  of  certain 
substances,  notably  manganese  dioxide.  By  mixing  such  a  substance 
with  the  chlorate  it  is  possible,  therefore,  to  expel  the  oxygen  rapidly 
at  a  lower  temperature  than  would  otherwise  have  to  be  employed.  The 
operation,  as  commonly  carried  out  in  the  laboratory,  is  as  follows : 

The  potassium  chlorate,  mixed  with  about  one  fourth  of  its  weight  of  man- 
ganese dioxide,  is  placed  in  a  suitable  vessel,  such  as  a  glass  flask,  which  is  pro- 
vided with  a  cork  and  glass  tube,  as  shown  in  A  (Fig.  6).   Upon  applying  a  gentle 
heat,  oxygen  is  evolved  and  passes  out  through 
the  tube  B.   It  is  evident  that  the  oxygen  at  first 
escaping  is  mixed  with  the  air  contained  in  the      ' 
flask.  In  a  short  time,  as  the  evolution  of  oxygen 
continues,  all  this  air  is  displaced,  and  the  pure 
oxygen  may  then  be  collected  by  bringing  the 
end  of  the  delivery  tube  under  the  mouth  of  a 
glass  cylinder   C,  which  has   been   filled  with  -p1G   7 

water  and  inverted  in  a  trough  of  water,  as  shown 

in  the  figure.  The  gas  rises  in  the  cylinder  and  displaces  the  water.  In  preparing 
larger  quantities  of  oxygen  a  copper  retort  (Fig.  7)  having  a  capacity  of  from  500 
to  1000  cc.  may  be  used  to  advantage  in  place  of  the  more  fragile  glass  flask. 

Although  the  complete  discussion  of  the  changes  which  take  place 
on  heating  potassium  chlorate  must  be  postponed  until  a  later  chapter, 


20 


GENERAL  CHEMISTRY 


it  is  possible  at  this  time  to  state  in  a  general  way  what  occurs.  The 
composition  of  potassium  chlorate  is  as  follows :  potassium,  31.9  per 
cent;  chlorine,  28.9  per  cent;  oxygen,  39.2  per  cent.  When  the  sub- 
stance is  heated,  a  series  of  changes  occur  which  finally  result  in  the 
liberation  of  all  the  oxygen  and  the  formation  of  a  white,  solid  com- 
pound of  potassium  and  chlorine,  known  as  potassium  chloride.  These 
facts  may  be  expressed  by  the  following  method,  in  which  the  names 
of  the  elements  present  in  each  compound  are  given  in  brackets  just 
below  the  name  of  the  compound : 

potassium  chlorate >•  potassium  chloride  -f-  oxygen 

rpotassiuml 

(potassium  | 
chlorine 

Lchlorme 
[oxygen      J 

As  to  the  way  in  which  the  manganese  dioxide  promotes  the  de- 
composition, it  may  be  said  at  once  that  we  do  not  know.  Apparently 
it  undergoes  no  change  during  the  reaction.  Certainly  it  contributes 
no  oxygen,  for  the  weight  of  the  latter  obtained  is  always  39.2  per 

cent  of  the  weight  of  the  chlorate 
used,  irrespective  of  the  presence  of 
manganese  dioxide.  This  is  but  one 
example  of  many  in  which  the  rate  of 
change  is  influenced  by  an  apparently 
inactive  substance.  Such  materials  are 
called  Catalytic  agents,  or  catalyzers^ 
and  we  shall  meet  with  them  fre- 
quently in  subsequent  pages. 

A  more  convenient,  although  more 
expensive,  method  for  preparing  oxy- 
gen in  the  laboratory  consists  in  adding  water  to  sodium  peroxide. 
These  two  substances,  when  brought  in  contact  with  each  other,  react 
in  such  a  way  as  to  liberate  oxygen.  At  the  same  time  there  is  formed 
a  compound,  consisting  of  sodium,  hydrogen,  and  oxygen,  known  as 
sodium  hydroxide : 

sodium  peroxide  4-  water >-  sodium  hydroxide  +  oxygen 

["sodium"!          ["hydrogen"!  podium     ~| 

[oxygen]          [oxygen  hydrogen 

[oxygen     J 

A  convenient  form  of  apparatus  for  generating  oxygen  by  this  method  is 
shown  in  Fig.  8.  The  peroxide  is  placed  in  the  flask  A  while  the  bulb  of  the 


FIG.  8 


OXYGEN  21 

separatory  funnel  B  is  filled  with  warm  water.  The  stopcock  C  is  then  turned 
until  the  water  enters  the  flask  drop  by  drop.  As  soon  as  the  water  comes  in 
contact  with  the  peroxide,  oxygen  is  evolved  which  escapes  through  D  into  the 
adjoining  bottle ;  after  bubbling  through  the  water  in  the  bottle,  it  passes  out 
and  is  collected  by  the  usual  method. 

Commercial  preparation.  A  number  of  methods  for  the  preparation 
of  oxygen  on  a  large  scale  have  been  employed  at  different  times.  In 
the  United  States,  at  the  present  time,  practically  all  of  the  oxygen 
prepared  commercially  is  obtained  either  from  liquid  air 
or  from  potassium  chlorate.  It  is  pumped  into  strong  steel 
tubes  under  great  pressure  (Fig.  9),  and  in  this  form  is  an 
article  of  commerce.  It  was  formerly  obtained  by  the  use 
of  barium  peroxide,  taking  advantage  of  the  reactions  already 
described.  This  process  was  known,  from  the  name  of  its_ 
inventor,  as  the  Brin  proc 

les.  Uncler  ordinary  conditions,  oxygen  is  a  color- 

AL  U"  alld  Under  orJmarv 

atmospheric  pressure  100  volumes  ot  water  dissolve  approx^- 
imately  4  yolumes  of  the  gas,  so  that  it  is  not  very  soluble 
and  may  be  collected  over  water  witri  little  16ssT~Une  liter 
of  oxygen,  measured  at  0°  and  under  a  pressure  equivalent 
to  that  of  a  column  of  mercury  760  mm.  in  height  (the  normal 
barometric  pressure  at  the  sea  level),  weighs  1.4290  g.  Since 
1 1.  of  air  under  the  same  conditions  weighs  1.2928  g.,  it  will 
be  seen  that  oxygen  is  1.105  times  as  heavy  as  air. 

Through  the  combined  effect  of  pressure  and  low  tempera- 
ture, oxygen  may  be  obtained  in  the  form  of  a  liquid.    To 
accomplish   the  liquefaction    the   temperature    must   be    at      FlG  9 
least  as  low  as  —118°,  at  which  temperature  a  pressure  of 
50  atmospheres  is  required  (see  critical  temperature,  p.  76).    At  still 
lower  temperatures  less  pressure  will  suffice.    Liquid  oxygen  has  a 
slightly  bluish  color  and  boils  at  —182.9°  under  a  pressure  of  1  atmos- 
phere.  It  is  strongly  attracted  by  a  magnet.  By  subjecting  this  liquid 
to  an  extremely  low  temperature,  Dewar,  an  English  investigator, 
succeeded  in  freezing  it  to  a  snowlike  solid  which  melts  at  —  235°. 

Chemical  conduct.  At  ordinary  temperatures  oxygen  is  only  a  mod- 
erately active  element  —  a  fact  which  may  be  inferred  from  our  experi- 
ence that  very  few  of  the  materials^  coming  under  common  observation 
are  acted  upon  by  the  oxygen  of  the  air  with  any  noticeable  rapidity. 


22 


GENERAL  CHEMISTRY 


FIG. 10 


With  rise  of  temperature  it  becomes  very  much  more  active.  At  ordi- 
nary temperatures,  for  example,  the  elements  sulfur,  iron,  and  carbon 
are  not  noticeably  acted  upon  by  oxygen,  while  in  the  case  of  phos- 
phorus, the  action,  though  slow,  is  quite  apparent.  If  now  the  temper- 
ature of  each  of  these  elements  is  slowly  raised,  the  reaction  becomes 
more  marked  and  the  phosphorus  soon  bursts  into  flame.  At  higher 
temperatures  the  sulfur,  carbon,  and  iron  likewise 
are  ignited.  If  each  of  the  elements,  as  soon  as 
ignited  in  the  air,  is  introduced  into  a  vessel  of 
pure  oxygen  (Fig.  10),  the  action  becomes  much 
more  energetic.  The  pale  blue  flame  of  the  burn- 
ing sulfur  is  greatly  increased  in  size  and  bright- 
ness, the  iron  throws  off  countless  sparks,  while 
the  phosphorus  and  carbon  burn  with  dazzling 
brilliancy.  In  each  case  the  action  increases  in 
intensity  as  it  progresses.  Many  compounds  act 
in  the  same  general  way.  Thus  wood,  coal,  oil, 
fats,  and  natural  and  artificial  gas  all  burn  readily 
in  air  and  more  brilliantly  in  pure  oxygen.  Indeed, 
there  are  but  few  elements  which,  under  the  proper  conditions  of 
temperature,  will  act  upon  so  many  other  elements  and  compounds  as 
does  oxygen. 

This  general  conduct  suggests  a  great  many  questions.  What  be- 
comes of  the  materials  when  they  burn  ?  Why  is  there  a  difference 
in  the  ease  of  ignition  ?  Why  do  not  all  substances  burn  ?  Why  is 
the  action  more  intense  in  oxygen  than  in  the  air,  and  why  does  it  be- 
come more  energetic  and  brilliant  as  it  progresses  ?  To  some  of  these 
questions  we  can  find  an  answer  at  once ;  others  will  occur  again 
many  times  in  our  study  and  in  the  end  will  remain  only  partially 
answered. 

The  nature  of  the  action  of  oxygen  upon  substances  ;  oxidation  ; 
oxidizing  agent.  By  means  of  experiments  it  is  possible  to  show  that 
tKe"action  of  oxygen  upon  another  element  consists  in  the  union  of 
the  two  elements  to  form  a  compound.  Thus,  when  sulfur  burns  in 
oxygen,  both  sulfur  and  oxygen  disappear  as  such,  and  in  their  place 
we  find  a  gaseous  compound  composed  of  the  two  elements.  Like- 
wise, when  phosphorus,  iron,  ajidcnrbrm  burn  in  oxygen,  there  are 
formed  compounds  of  these  elements  with  oxygen.  The  action  of 
oxygen  upon  compounds  is  similar  to  its  action  upon  elements,  and 


OXYGEN  23 

consists  in  the  union  of  the  oxygen  with  one  or  more  of  the  elements 
present  in  the  compound,  or,  in  some  cases,  with  the  compound  as  a 
whole.  Thus,  when  the  gaseous  compound  of  hydrogen  and  sulfur 
known  as  hydrogen  sulfide  burns  in  a  limited  supply  of  air,  only  the 
hydrogen  present  combines  with  the  oxygen,  while  sulfur  is  liberated. 
If  the  supply  of  air  is  sufficient,  however,  to  furnish  the  necessary 
oxygen,  then  both  the  hydrogen  and  the  sulfur  present  combine  with 
oxygen. 

Tbp-  gftperal  term  oxidation  is  applied  to  all  such  processes  as  those 
Described  abovej.  in  which  any  substance  or  its  constituent  parts  com- 
bines with  oxygen.  Thus  we  speak  of  the  oxidation  of  phosphorus  or 
sulfur  by  the  air  or  by  pure  oxygen,  and  we  say  that  these  elements 
readily  undergo  oxidation.  The  material  which  supplies  the  oxygen 
is  called  the  oxidizing  agent.  In  the  examples  just  mentioned  the  air 
or  pure  oxygen  itself  is  the  oxidizing  agent,  but  in  many  cases  the 
oxygen  is  supplied  by  some  compound  such  as  potassium  chlorate  or 
sodium  peroxide. 

Oxides ;  products  of  oxidation.  When  any  element  combines  with 
oxygen,  the  reyiltin|y  compound  is  known  as  an  oxide  of  that  element. 
Thus  the  compound  formed  by  the  union  of  sulfur  with  oxygen  is 
known  as  an  oxide  of  sulfur.  Likewise,  when  phosphorus,  iron,  and 
carbon  combine  with  oxygen,  the  resulting  compounds  formed  are 
oxides.  The  particular  oxide  or  oxides  formed  in  the  oxidation  of  any 
substance  are  known  in  general  as  the  products  of  oxidation  of  that 
substance. 

Oxides  of  nearly  all  the  elements  have  been  prepared,  the  only 
exceptions  being  that  of  fluorine  and  of  a  group  of  gaseous  ele- 
ments occurring  chiefly  in  the  air  —  helium,  neon,  argon,  krypton, 
xenon.  Moreover,  some  of  the  elements  combine  with  oxygen  in  sev- 
eral different  proportions,  so  that  we  may  have  more  than  one  oxide 
of  the  same  element.  Barium  forms  two  such  compounds,  barium 
oxide  and  barium  peroxide,  to  which  reference  has  already  been  made 
(p.  18)..  It  is  evident,  therefore,  that  the  oxides  constitute  a  large  and 
important  class  of  compounds. 

Some  of  the  oxides  are  invisible  gases,  as  is  true  of  the  oxide  of 
sulfur  and  of  carbon.  In  a  few  cases  the  oxide  is  a  liquid,  the  most 
familiar  example  being  water,  which  is  an  oxide  of  hydrogen.  In  the 
great  majority  of  cases,  however,  the  oxides  are  solids,  which  is  true 
of  those  of  iron  and  phosphorus.  It  is  easy  to  understand,  therefore, 


24 


GENERAL   CHEMISTRY 


why  such  elements  as  sulfur  and  carbon  completely  vanish  on  burn- 
ing, leaving  no  ash,  while  other  elements,  such  as  iron  and  phosphorus, 
leave  a  solid  residue. 

Weight  relations  in  oxidation.  If  it  is  true  that  oxidation  is  essen- 
tially the  union  of  oxygen  with  other  elements  or  compounds,  it  must 
necessarily  follow  that  the  weight  of  the  product  formed  is  greater 
than  that  of  the  substance  oxidized  or  burned,  although  our  common 
experience  with  fuels  would  hardly  lead  us  to  such  a  conclusion. 
That  the  weight  does  increase  may  be  demonstrated  by  arranging 
an  experiment  in  which  the  oxidation  takes  place  on  one  pan  of  a 
balance.  For  example,  some  powdered  iron  may  be  placed  in  a  small 
porcelain  dish  and  accurately  counterpoised  on  the  balance.  The  iron 
may  then  be  ignited  by  directing  a  hot  flame  upon  it.  As  the  oxi- 
dation proceeds,  it  will  be  seen  that  the  scale  pan  supporting  the  dish 

sinks.  A  modification  of  this  experi- 
ment will  show  a  similar  gain  in  weight 
during  the  burning  of  a  candle. 

The  candle  is  arranged  on  a  balance 
pan  shown  in  Fig.  11.  Over  it  is  suspended 
a  wide  glass  tube  (a  lamp  chimney)  loosely 
filled  with  sticks  of  sodium  hydroxide,  a 
substance  which  will  absorb  both  water  and 
oxide  of  carbon.  When  the  whole  appa- 
ratus has  been  brought  to  an  equilibrium 
on  the  balance,  the  candle  is  lighted.  As 
it  burns,  the  materials  of  which  it  is  com- 
posed (chiefly  hydrogen  and  carbon)  are 
oxidized,  forming  oxide  of  hydrogen  (water) 

and  of  carbon,  which  are  drawn  up  through  the  chimney  and  absorbed  by  the 
sodium  hydroxide.  The  balance  pan  sinks  as  the  oxidation  progresses,  indicat- 
ing an  increase  in  weight. 

Such  experiments  as  these  have  been  carried  out  with  the  greatest 
care  by  skillful  workers,  and  it  has  been  shown  that  the  weight  of  the 
material  oxidized,  plus  that  of  the  oxygen  used  up,  exactly  equals  the 
weight  of  the  products  of  the  oxidation. 

Heat,  temperature,  and  light  in  oxidation.  In  the  burning  of  the 
elements  we  have  been  considering,  as  well  as  of  such  familiar  fuels  as 
wood,  coal,  and  oil,  it  is  noticed  that  at  first  the  heat  given  off 
does  not  seem  to  be  very  great,  and  the  light  is  quite  feeble.  As  the 
oxidation  proceeds,  the  temperature  rises  quite  rapidly  and  the  light 
grows  brighter,  in  some  cases  becoming  almost  blinding.  This  conduct 


FIG. 11 


OXYGEN  25 

is  readily  understood  if  we  keep  clearly  before  us  the  relation  between 
heat,  temperature,  and  light. 

By  heat  we  refer  to  a  definite  quantity  .of  one  form  of  energy, 
namely,  heat  energy,  and  we  measure  this  in  calories  (p.  7).  By 
Jemperati,ire  wejnean  the  intensity  of  this  p"p.rpy.  and  this  WP  mga«- 
_urjEL_by  a  jjiermometer.  An  illustration  will  make  the  distinction 
clearer.  Let  us  suppose  a  given  weightof  some  material  to  contain 
100  cal.  of  heat.  If  now  heat  is  applied  to  this  material  until  an- 
other 100  cal.  is  added  to  it,  the  weight  of  the  material  remains 
constant,  while  the  quantity  of  heat  is  doubled.  The  quantity  of 
heat  associated  with  a  given  weight  of  matter  has  been  increased,  and 
we  say  that  the  intensity  is  greater,  this  being  indicated  by  a  higher 
temperature.  If,  however,  we  add  to  the  material  containing  100  cal. 
of  heat  an  equal  weight  of  the  same  material,  also  containing  100  cal., 
the  quantity  of  heat  is  doubled,  but  the  quantity  of  matter  is  doubled 
at  the  same  time.  In  this  case  the  intensity  remains  unchanged,  and 
the  temperature  is  the  same  as  before. 

Accordingly,  when  chemical  energy  is  changed  into  heat  in  the 
process  of  oxidation,  under  conditions  such  that  the  loss  through 
radiation  and  conduction  is  small,  all  the  materials  concerned  in  the 
reaction  have  large  additions  to  their  heat  energy,  and  their  tempera- 
ture consequently  rises.  At  high  temperatures  (depending  on  the 
nature  of  the  materials)  a  part  of  this  heat  is  changed  into  light 
energy,  the  transformation  being  greater  the  higher  the  temperature. 
Increasing  brightness  is  therefore  an  indication  of  rising  temperature. 

Reactions  in  which  heat  is  set  free,  like  those  we  have  been  con- 
sidering, are  called  e.rothermic.  It  is  evident  that  in  such  reactions  the 
termoerature  will  begin  to  rise  just  as  soon  as  the  rate  at  which  heat 
is  liberated  exceeds  that  at  which  it  is  lost  through  radiation  and  con- 
duction. The  reaction  will  then  proceed  without  further  application 
of  external  heat  and  with  increasing  rapidity. 

In  many  reactions  heat  is  absorbed  and  must  be  supplied  from 
external  sources  if  the  action  is  to  continue.  Reactions  of  this  kind 
are  called.,  endothermic.  The  decomposition  of  potassium  chlorate  and 
mercuric_oxide  belong  to  this  class.  Evidently  such  reactions  can 
never  become  self-supporting. 

Combustion.  When  a  reaction  proceeds  rapidly  enough  to  produce 
light,  it  is  called  combustion.  Naturally  such  reactions  are  exothermic 
in  character.  The  most  familiar  examples  of  combustion  are  those  in 


26  GENERAL  CHEMISTRY 

which  substances  burn  in  air  or  oxygen,  and  which  are  therefore  also 
oxidations.  We  shall  meet  with  cases  of  combustion  in  subsequent 
pages,  however,  in  which  there  is  no  oxygen  taking  part  in  the 
reaction.  Ordinarily,  however,  when  we  speak  of  a  combustible  sub- 
stance, we  mean  one  that  will  burn  in  air  or  oxygen. 

Incombustible  substances.  The  question  naturally  rises,  Why  do  not 
all  substances  burn  ?  In  the  case  of  many  substances,  such  as  the  com- 
pounds formed  in  combustion,  the  answer  is  very  evident,  for  they 
already  contain  all  the  oxygen  with  which  they  are  capable  of 
uniting,  at  least  under  the  conditions  which  prevail  during  combus- 
tion. Many  oxides,  such  as  water,  are  of  this  class,  as  well  as  most  of 
the  substances  which  constitute  the  solid  crust  of  the  earth.  The 
materials  of  which  fireproof  buildings  are  made  —  brick,  tile,  cement, 
plaster,  asbestos  —  are  also  of  this  character.  Some  of  the  metals  like 
iron  will  burn,  but  only  at  such  high  temperatures  that  they  are 
practically  incombustible  under  all  ordinary  conditions.  Other  sub- 
stances, however,  such  as  the  elements  fluorine  and  argon,  do  not 
combine  with  oxygen  under  any  known  condition. 

Spontaneous  combustion.  Our  attention  is  frequently  called  to  the 
fact  that  certain  substances  sometimes  take  fire  spontaneously,  that  is, 
without  the  application  of  any  external  heat.  For  example,  a  piece 
of  phosphorus  exposed  to  the  air  in  a  moderately  warm  room  will 
soon  burst  into  flame.  Such  substances  are  always  easily  oxidized, 
and  much  heat  is  set  free  in  the  process.  In  contact  with  air  a  slow 
oxidation  takes  place  at  ordinary  temperatures,  and  if  the  heat  radia- 
tion is  not  too  great,  the  temperature  rises  and  the  reaction  is  hastened, 
with  the  result  that  ignition  presently  occurs.  Certain  oils,  such  as 
the  common  linseed  oil  used  in  paints,  readily  undergo  oxidation,  and 
oily  rags  carelessly  left  by  painters  often  cause  disastrous  fires.  In  a 
similar  way  coal,  especially  when  stored  in  large  quantities  and  in 
a  warm  place,  as  in  the  hold  of  a  ship,  may  take  fire  spontaneously. 

Slow  oxidation.  More  frequently  the  heat  set  free  when  substances 
slowly  oxidize  in  the  air  is  conducted  away  about  as  fast  as  it  is  liber- 
ated, so  that  the  temperature  does  not  greatly  change.  These  slow 
oxidations  are  often  of  great  economic  importance.  Most  of  the  metals 
slowly  rust  in  the  air,  and  in  some  cases  this  process  is  simply  one  of 
oxidation,  the  product  being  an  oxide  of  the  metal.  In  other  cases, 
as  with  iron,  the  reaction  is  more  complex,  since  water  takes  part  in 
it  and  the  rust  contains  hydrogen  as  well  as  oxygen  and  the  metal. 


OXYGEN  27 

Many  of  the  changes  involved  in  the  various  processes  of  growth 
and  decay  of  organisms  belong  to  this  same  class  of  slow  oxidations. 
The  oxidation  of  the  food  which  we  eat  keeps  the  body  at  the  most 
efficient  temperature.  Broken-down  tissue  is  changed  by  oxidation 
into  forms  in  which  it  can  be  easily  eliminated  from  the  body.  Much 
of  the  refuse  matter  from  organisms,  which  would  soon  become  a 
nuisance  if  not  a  danger  to  health,  is  oxidized  for  the  most  part  into 
water  and  oxide  of  carbon,  both  of  which  are  again  absorbed  by 
growing  plants.  It  is  a  very  wonderful  adjustment  by  which  all 
these  processes  work  in  harmony  with  each  other. 

Speed  of  oxidation.  The  description  of  the  conditions  which  lead  to 
spontaneous  combustion  has  already  made  it  clear  that  a  given  oxida- 
tion may  take  place  at  very  different  rates.  It  is  rather  natural  to 
infer  that  when  the  oxidation  of  a  definite  weight  of  any  substance 
proceeds  very  slowly,  less  heat  is  evolved  than  when  it  proceeds  rapidly, 
for  in  the  latter  case  a  much  higher  temperature  is  reached.  Very 
exact  measurements  have  shown  that  such  an  inference  is  wrong ;  the 
total  heat*  evolved  in  the  two  cases  is  precisely  the  same  irrespective  of 
the  rate,  provided  only  that  the  same  products  and  the  same  weight 
of  them  are  formed.  It  is  also  true  that  in  general  the  same  products 
are  formed,  though  this  is  not  always  so. 

It  is  a  matter  of  much  interest  for  us  to  inquire  what  conditions  in- 
fluence the  rate  or  speed  of  a  reaction,  the  speed  being  measured  by  the 
weight  of  material  undergoing  change  in  a  unit  of  time.  Such  knowl- 
edge will  have  much  practical  value,  for  in  various  industrial  proc- 
esses increase  of  speed  of  reaction  means  great  saving  of  time.  In  the 
purchase  of  fuel  it  is  the  heating  value  we  really  pay  for,  and  some- 
times we  wish  to  use  this  so  as  to  get  as  much  heat  as  possible  in  a 
short  time,  while  for  other  purposes  we  m.ay  desire  to  obtain  the  heat 
slowly  but  evenly. 

Conditions  which  influence  the  speed  of  oxidation.  For  the  present  we 
need  consider  only  three  factors  which  affect  the  speed  of  oxidation. 

1.  Effect  of  temperature.  The  most  obvious  influence  affecting  the 
speed  of  oxidation  is  temperature.  At  a  high  temperature  oxidation 
takes  place  rapidly;  as  the  temperature  is  lowered  the  speed  de- 
creases, until  at  ordinary  temperatures  it  may  be  impossible  to  detect 
any  action  whatever.  It  is  exceedingly  improbable,  however,  that  the 
process  entirely  ceases ;  it  merely  becomes  so  slow  that  we  have  no 
means  of  detecting  any  change. 


28  GENERAL  CHEMISTRY 

2.  Effect  of  concentration.   It  is  evident  that  anything  which  increases 
the  quantity  of  oxygen  in  contact  with  the  surface  of  the  burning  sub- 
stance will  tend  to  hasten  the  reaction.    One  reason  why  substances 
burn  more  rapidly  in  oxygen  than  in  air  is  that  the  latter  is  only  one 
fifth  oxygen,  so  that  a  body  burning  in  air  is  at  one  time  in  contact 
with  only  one  fifth  as  much  oxygen  as  when  it  is  burning  in  the  pure 
gas.    Instead  of  increasing  the  concentration  of  the  oxygen  we  may 
often  hasten  the  oxidation  by  extending  the  surface  of  the  solid  sub- 
stance.   A  log  of  wood  burns  more  slowly  than  the  same  wood  split 
into  kindling.    A  lump  of  coal  burns  rather  slowly,  but  when  finely 
powdered  and  suspended  in  the  air  as  dust,  it  burns  almost  instan- 
taneously, and  with  explosive  violence. 

3.  Catalysis.    We  have  already  seen  that  potassium  chlorate  gives 
up  its  oxygen  much  more  rapidly  under  the  catalytic  influence  of 
manganese  dioxide  than  when  heated  alone.    In  a  similar  way  the 
speed  of  oxidation  may  sometimes   be   increased  by  the   action   of 
some  suitable  catalytic  agent.    Thus,  when  the  oxide  formed  in  the 
combustion  of  sulfur  (called  sulfur   dioxide)  is   heated  with  oxy- 
gen under  the  proper  conditions,  it  slowly  takes  up  an  additional 
quantity  of  oxygen  to  form  a  new  compound  known  as  sulfur  triox- 
ide.    It  has  been  found  that  the  speed  of  this  oxidation  may  be  greatly 
increased  by  the  presence  of  certain  catalytic  agents,  such  as  finely 
divided  platinum,  and  this  discovery  has  led  to  marked  improvements 
in  the  manufacture  of  sulfuric  acid.    In  like  manner  the  presence  of 
a  trace  of  moisture  greatly  increases  the  speed  of  oxidation  of  many 
substances. 

Importance  of  oxygen.  The  great  importance  of  oxygen  in  nature 
is  evident  from  the  facts  which  have  already  been  presented  in  this 
chapter.  It.  is  a  constituent  of  the  great  majority  of  the  compounds 
which  collectively  constitute  the  solid  earth,  the  living  creatures  upon 
it,  and  the  water  which  covers  so  much  of  its  surface,  while  the  atmos- 
phere is  a  great  reservoir  from  which  a  supply  of  the  free  element  can 
be  drawn  at  any  time. 

Free  oxygen  is  essential  to  the  life  of  all  organisms,  with  the 
exception  of  some  of  the  lowest  forms.  Aquatic  animals  obtain  the 
necessary  oxygen  from  the  air  dissolved  in  the  water  in  which  they 
live.  Free  oxygen  also  plays  a  prominent  part  in  the  decomposi- 
tion of  refuse  organic  matter,  much  of  it  being  oxidized  into  harm- 
less gases.  It  is  noteworthy,  however,  that  the  oxidation  of  such 


OXYGEN  29 

matter  takes  place  only  in  the  presence  of  certain  minute  forms  of 
living  organisms  known  as  bacteria. 

Free  oxygen  is  also  utilized  in  a  great  variety  of  industrial  processes, 
but  for  most  of  these  air  answers  every  purpose,  since  the  nitrogen 
which  it  contains  does  not  seriously  interfere.  Pure  oxygen  finds  ap- 
plication in  quite  a  variety  of  scientific  experiments,  in  the  production 
of  very  high  temperatures,  and  in  the  treatment  of  certain  diseases 
in  which  the  patient  is  unable  to  inhale  sufficient  air  to  supply  the 
necessary  quantity  of  oxygen. 

The  definiteness  of  chemical  processes.  Throughout  this  chapter  at- 
tention has  been  repeatedly  directed  to  the  fact  that  chemical  processes 
involve  definite  weights  of  matter.  For  example,  the  composition  of 
a  number  of  compounds  has  been  expressed  in  exact  percentages, 
since  experiment  has  shown  that  these  always  have  precisely  the 
composition  stated,  irrespective  of  the  source  from  which  they  are  ob- 
tained or  the  method  by  which  they  are  prepared.  After  extensive 
investigation  of  a  very  large  number  of  compounds,  chemists  have 
concluded  that  this  constancy  of  composition  is  a  characteristic  of 
every  true  compound,  and  a  statement  of  this  characteristic  is  com- 
monly called  the  law  of  definite  composition. 

In  like  manner  the  chemical  changes  which  compounds  undergo 
are  always  perfectly  definite  under  stated  conditions.  Thus,  when 
potassium  chlorate  is  heated,  for  every  100  g.  decomposed  there  re- 
sult 39.2  g.  of  oxygen  and  60.8  g.  of  potassium  chloride.  When  iron 
burns  in  oxygen,  100  g.  of  iron  combines  with  38.20  g.  of  oxygen  to 
form  138.20  g.  of  oxide  of  iron.  If  less  than  38.20  g.  of  oxygen  is 
present,  then  a  corresponding  amount  of  iron  will  remain  unchanged. 
On  the  other  hand,  if  more  than  38.20  g.  of  oxygen  is  present,  then 
all  the  iron  will  be  changed  into  the  oxide,  and  the  excess  of  oxygen 
will  remain  unaltered.  The  actual  experiments  which  justify  these 
conclusions  will  come  before  us  from  time  to  time  as  we  proceed. 

The  phlogiston  theory  of  combustion.  Before  leaving  the  topic  of  combustion  it 
will  be  of  interest  to  contrast  our  present  ideas  with  those  of  the  chemists  of  a 
few  centuries  ago,  or  of  the  alchemists,  as  they  were  sometimes  termed.  Of  the 
many  conceptions  which  have  been  held  at  different  times,  that  which  is  known 
as  the  phlogiston  theory  had  by  far  the  greatest  influence  upon  the  develop- 
ment of  chemistry.  This  theory  was  advanced  by  Becher  (1635-1682)  and  was 
greatly  extended  and  developed  by  the  distinguished  German  professor  Stahl 
(1660-1734).  According  to  this  theory  every  combustible  substance  contains 
more  or  less  of  a  material,  or  "  principle,"  called  phlogiston,  the  escape  of  which 


30 


GENERAL  CHEMISTRY 


constitutes  combustion.    The  ash  remaining  represents  the  original  substance 
minus  phlogiston.   Substances  which  leave  no  ash  are  nearly  pure  phlogiston. 

When  it  is  remembered  that  at  that  time  gases  were  little  understood,  oxygen 
unknown,  and  heat  and  light  regarded  as  material  substances  given  off  during  com- 
bustion, and  that  in  this  process  something  very  evidently  did  escape  (namely,  the 
gaseous  oxides),  it  will  be  seen  that  the  theory  was  a  reasonable  one.  It  was 
known  that  metals  increased  in  weight  during  combustion,  but  little  importance 
was  at  that  time  attached  to  weight  relations,  and  some  adherents  of  the  theory 
even  assumed  that  phlogiston  had  negative  weight,  that  is,  weighed  less  than 
nothing.  It  was  also  difficult  to  explain  why  combustion  required  the  presence 
of  air.  The  theory  was  almost  universally  held  for  a  hundred  years,  and  was  given 
up  only  after  oxygen  had  been  discovered  and  Lavoisier  had  demonstrated  the 
true  nature  of  combustion.  Even  then  some  chemists  did  not  accept  the  new 
ideas.  Priestley,  whose  discovery  of  oxygen  contributed  so  much  to  the  true 

explanation  of  combustion,  never 
gave  up  the  phlogiston  theory.  In 
fact,  one  of  his  last  publications 
was  entitled  "The  Doctrine  of 
Phlogiston  Established." 

The  demonstration  of  the  true 
nature  of  combustion  must  be 
regarded  as  one  of  the  greatest 
achievements  in  the  history  of 
chemical  science,  for  our  modern 
views  date  from  it.  It  is  therefore 
worth  while  to  learn  something 
of  the  methods  which  Lavoisier 
employed.  Fig.  12  shows  the  form 
of  apparatus  used  in  one  of  his  most  important  experiments.  A  bell  jar  A  full 
of  air  was  placed  on  a  basin  of  mercury  B,  and  a  retort  was  arranged  in  such 
a  way  that  its  delivery  tube  passed  under  the  edge  of  the  bell  jar  and  communi- 
cated with  the  included  air.  Lavoisier  placed  4  oz.  of  mercury  in  the  retort  and 
heated  it  for  12  days.  As  the  heating  progressed,  he  observed  that  the  mercury 
was  gradually  converted  into  a  red  solid  (oxide  of  mercury)  and  that  this  change 
was  attended  by  a.  contraction  in  the  volume  of  the  air  in  the  bell  jar,  as  shown 
by  the  rise  of  mercury  in  it.  The  contraction  amounted  to  from  7  to  8  cu.  in. 
He  then  placed  the  red  solid  in  a  small  retort  and  decomposed  it  by  heating  it  to 
a  higher  temperature,  obtaining  the  original  mercury  once  more,  together  with 
from  7  to  8  cu.  in.  of  oxygen.  He  also  determined  the  weight  relations  of  the 
several  substances  taking  part  in  these  changes,  and  was  able  to  prove  that  when 
tin  is  heated  in  contact  with  the  air,  it  combines  with  oxygen,  and  that  the 
increase  in  weight  is  equal  to  that  of  the  oxygen  absorbed.  In  addition  to  these 
experiments,  Lavoisier  burnt  phosphorus  in  a  limited  amount  of  air  confined 
over  mercury  and  found  that  only  a  portion  of  the  phosphorus  burned,  the 
volume  of  the  air  at  the  same  time  being  diminished  about  one  fifth.  He  ob- 
served that  a  white  solid  was  formed  when  the  phosphorus  burned  and  proved 
this  to  be  an  oxide  of  phosphorus.  These  experiments  were  all  carried  out  with 
an  exactness  remarkable  for  that  period. 


FIG. 12 


OXYGEN 


31 


OZONE 

Historical.  In  1785  the  Dutch  chemist  Van  Maram  observed  that 
oxygen  through  which  electric  sparks  have  been  passed  acquires  a 
peculiar  odor.  No  further  attention  was  given  to  the  observation 
until  1840,  when  Schonbein  showed  that  the  odor  observed  by  Van 
Marum  was  due  to  the  presence  of  a  new  gaseous  substance  formed 
from  oxygen  by  the  action  of  the  electric  discharge.  Schonbein  called 
the  new  substance  ozone,  the  name  being  derived  from  a  Greek  word 
signifying  "  to  smell."  Finally,  in  1856,  Andrews  proved  that  ozone  is 
elementary  in  character,  since  no  material  substance  is  added  to  oxygen 
in  its  conversion  into  ozone. 

Preparation.  The  method  used  in  preparing  ozone  is  the  same  in 
principle  as  that  which  led  to  the  discovery  of  the  substance,  and 
consists  in  subjecting  oxygen  or  air  to  the  influence  of  an  electric 
discharge.  It  has  been  found  that  the  yield  of 
ozone  can  be  greatly  increased  by  using  the 
silent  electric  discharge  rather  than  electric 
sparks,  since  in  this  way  it  is  possible  not  only 
to  bring  a  larger  volume  of  oxygen  under  the 
influence  of  the  discharge,  but  to  do  this  with- 
out raising  the  temperature  to  any  marked 
degree,  and  thus  to  avoid  decomposition  of  the 
ozone  by  the  heat.  An  apparatus  used  for  pre- 
paring ozone  by  this  method  is  termed  an  ozon- 
izer  and  consists  essentially  of  two  conducting 
metallic  surfaces,  separated  by  a  dielectric  such 
as  glass  or  mica,  and  so  constructed  that  a 
current  of  air  may  be  passed  between  the  two 
conducting  surfaces. 

Fig.  13  represents  a  common  form  of  ozonizer.  Oxygen  or  air  enters  at  A  and 
follows  the  course  indicated  by  the  arrows.  The  conducting  surfaces  B  and  C 
are  separated  by  a  glass  dielectric  D.  Wires  leading  from  an  induction  coil  are 
connected  with  B  and  C.  As  the  oxygen  passes  upward  between  the  conducting 
surfaces  it  is  subjected  to  the  influence  of  the  electric  discharge,  and  a  portion 
of  the  element  is  thereby  changed  into  ozone. 

To  obtain  the  greatest  yield,  the  oxygen  should  be  cold  and  free 
from  moisture.  Under  ordinary  conditions  only  a  very  small  percent- 
age of  the  oxygen  is  transformed  into  ozone.  Recently,  however, 
Harries  has  obtained  a  yield  of  from  18  to  19  per  cent  by  means  of  an 


D 

3* 

\ 

I 

-c 

1 
t 

V 

/ 

B    B 

1 

/ 

X_ 

°t 
C 

i 
I 

> 

J- 

—  1_ 

i 

= 

~ 

—  ^r 

1 

ETJ 

nrz 

z.— 

— 

FIG. 13 


32 


GENERAL  CHEMISTRY 


improved  ozonizer.  To  prepare  pure  ozone  the  mixture  of  oxygen  and 
ozone  is  cooled  with  liquid  air.  The  ozone,  being  much  more  readily 
condensed  than  oxygen,  is  thereby  obtained  in  a  liquid  state  almost 
free  from  oxygen.  By  evaporating  this  liquid  and  again  condensing 
the  gas  the  oxygen  present  is  gradually  separated  and  pure  ozone 
thus  obtained. 

Ozone  is  also  formed  in  reactions  in  which  oxygen  is  liberated  at 
low  temperatures,  as  in  the  decomposition  of  water  by  the  electric 
current.    It  is  also  produced  in  some  oxidations,  as 
when  moist  phosphorus  slowly  oxidizes  in  air. 

Its  formation  by  the  oxidation  of  phosphorus  may  be 
shown  by  partially  covering  with  water  a  few  pieces  of 
stick  phosphorus  placed  in  the  bottom  of  a  jar  (Fig.  14). 
The  presence  of  ozone  in  the  air  in  the  jar  is  soon  indi- 
cated by  its  characteristic  odor,  as  well  as  by  the  property 
it  possesses  of  imparting  a  blue  color  to  strips  of  paper 
(yl)  previously  dipped  into  a  solution  of  potassium  iodide 
and  starch.  The  ozone  acts  upon  the  potassium  iodide, 
liberating  the  iodine  present,  which  in  turn  forms  with 
the  starch  a  blue  substance. 


FIG. 14 


It  is  of  much  importance  to  note  that  the  transformation  of  oxygen 
into  ozone  is  accompanied  by  a  change  in  volume,  3  volumes  of  oxygen 
forming  2  volumes  of  ozone. 

The  preparation  of  ozone  from  oxygen  is  a  reversible  change.  That 
only  a  relatively  small  percentage  of  the  oxygen  subjected  to  an  elec- 
tric discharge  is  transformed  into  ozone  is  due  to  the  fact  that  under 
all  ordinary  conditions  ozone  changes  back  into  oxygen.  These  two 
opposite  changes  may  be  represented  as  follows,  the  arrows  indicating 
the  direction  in  which  each  change  proceeds : 


oxygen 


ozone 


When  the  speed  with  which  the  oxygen  changes  into  ozone  exactly 
equals  the  speed  with  which  the  resulting  ozone  changes  back  into 
oxygen,  the  two  changes  are  said  to  be  in  a  state  of  equilibrium.  It  is 
evident  that  the  percentage  of  ozone  present  when  equilibrium  is 
reached  represents  the  maximum  yield  under  the  conditions  of  the  ex- 
periment. Such  reactions  as  the  above  are  termed  reversible  reactions. 
Properties  and  conduct.  Ozone  is  a  gas  of  pale  blue  color  and  char- 
acteristic odor.  It  is  1.5  times  as  heavy  as  oxygen,  1  1.  of  the  gas 
weighing  2.144  g.  When  cooled  sufficiently  it  condenses  to  a  deep  blue 


OXYGEN  33 

liquid  which  boils  at  —  119°  and  is,  like  liquid  oxygen,  strongly  mag- 
netic. Ozone  resembles  oxygen  in  its  chemical  conduct,  but  is  much 
more  active  and  is  therefore  a  very  powerful  oxidizing  agent.  Such 
metals  as  silver  and  mercury,  which  are  not  easily  acted  upon  by  oxygen, 
quickly  tarnish  in  ah*  containing  ozone.  It  likewise  oxidizes  many 
organic  dyes  into  colorless  compounds,  acting  as  a  bleaching  agent. 
When  pure  it  is  a  dangerous  explosive,  owing  to  its  spontaneous  con- 
version into  oxygen,  in  which  change  675  cal.  of  heat  are  liberated  for 
each  gram  of  ozone  converted. 

In  view  of  its  great  activity  it  is  doubtful  whether  ozone  is  ever 
present  in  the  air  in  appreciable  quantities.  It  is  doubtless  formed  in 
a  number  of  natural  processes,  as  by  lightning  discharge,  but  it  must 
very  speedily  disappear  through  chemical  action  upon  oxidizable  mate- 
rials which  are  always  present. 

Uses  of  ozone.  Ozone  finds  increasing  commercial  applications,  all 
based  on  its  strong  oxidizing  properties.  Thus  it  is  used  as  a  bleach- 
ing agent,  as  a  disinfectant,  and  as  an  oxidizing  agent  in  the  prepara- 
tion of  a  number  of  useful  products.  It  is  also  being  used,  especially 
in  Europe,  for  the  sterilization  of  drinking  waters.  For  these  various 
uses  the  pure  substance  is  not  employed,  but  ozonized  air  which  is 
prepared  as  it  is  needed. 

The  difference  between  oxygen  and  ozone.  Experiment  has  shown 
that  it  is  possible  to  change  pure  oxygen  into  ozone  or  ozone  into 
oxygen  without  change  in  weight,  so  that  the  difference  between  these 
two  distinct  substances  cannot  lie  in  the  material  of  which  they  are 
composed.  The  reason  for  their  difference  must  therefore  be  sought 
in  the  energy  relations  which  exist  between  them. 

In  the  transformation  of  oxygen  into  ozone  by  the  usual  method, 
electrical  energy  is  used  up,  and  this  must  be  accounted  for,  since  it 
cannot  be  lost.  When  ozone  spontaneously  changes  into  oxygen,  675 
caL  of  heat  are  liberated  for  each  gram  so  changed,  and  when  it  acts 
as  an  oxidizing  agent,  the  heat  evolved  is  correspondingly  greater 
than  when  the  same  oxidation  is  accomplished  by  oxygen.  Mechan- 
ical work  also  is  concerned  in  the  transformation,  for  there  is  a  change 
in  the  volume  (p.  32)  and  a  corresponding  one  in  the  density  of 
the  gas.  The  simplest  way  in  which  to  regard  these  relations  is  to  as- 
sume that  in  the  transformation  of  oxygen  into  ozone  the  electrical 
and  mechanical  energy  absorbed  is  stored  up  as  chemical  energy  in  the 
ozone.  When  ozone  spontaneously  decomposes  or  acts  upon  other 


34  GENERAL  CHEMISTRY 

substances,  this  excess  energy  reappears  as  heat.  Likewise  when  ozone 
is  formed  by  chemical  reactions,  a  part  of  the  chemical  energy  of  the 
reacting  substances  is  transferred  to  the  ozone. 

It  should  be  noted  that  the  conversion  of  oxygen  into  ozone  is  a 
perfectly  definite  change,  just  as  with  any  other  chemical  reaction.  It 
is  not  possible  to  add  increasing  quantities  of  chemical  energy  to 
oxygen  until  we  finally  get  ozone,  but  a  definite  weight  of  oxygen, 
will  always  give  a  definite  weight  of  ozone,  with  no  intermediate  stages. 
We  shall  find,  as  we  progress  in  our  study,  that  many  other  substances, 
compounds  as  well  as  elements,  can  be  obtained  in  two  or  more  forms 
of  the  same  percentage  composition  but  very  different  in  their  energy 
content. 


CHAPTER  III 

HYDROGEN 

Historical.  So  far  as  is  known,  hydrogen  was  first  prepared  by 
^Paracelsus  (1493-1541),  but  the  English  investigator  Cavendish  is 
commonly  regarded  as  its  discoverer,  since  he  was  the  first  to  obtain 
it  in  pure  condition  and  to  recognize  it  as  an  independent  substance 
different  from  other  known  inflammable  gases.  This  was  in  1766. 
Cavendish  termed  the  gas  inflammable  air,  but  later,  when  it  was 
found  that  it  was  a  constituent  of  water,  Lavoisier  renamed  it  hydrogen, 
signifying  "water  former." 

Occurrence.  Hydrogen  in  the  free  condition  is  sometimes  found  in 
gases  issuing  from  the  earth  in  volcanic  regions.  It  has  also  been 
found  in  samples  of  air  collected  in  different  localities,  but  only  in 
very  minute  quantities.  The  spectroscope  reveals  large  quantities  of 
the  free  element  in  the  gases  surrounding  the,  sun  and  certain  other 
stars.  Combined  with  oxygen  in  the  form  of  water  it  is  widely 
distributed.  Combined  with  carbon  it  forms  a  large  number  of 
compounds  known  as  hydrocarbons,  which  constitute  by  far  the 
greater  part  of  ordinary  natural  gas  and  petroleum.  It  is  likewise 
a  constituent  of  the  compounds  present  in  living  organisms  and 
of  most  of  the  products  derived  from  them,  such  as  sugar,  starch, 
and  butter.  It  is  invariably  a  constituent  of  the  compounds  known 
as  acids. 

Preparation  of  hydrogen.  Since  hydrogen  does  not  naturally  occur 
in  the  free  condition  to  any  extent,  it  must  be  prepared  by  liberating 
it  from •  its  compounds.  The  compounds  most  commonly  used  as 
sources  of  hydrogen  are  (1)  water,  (2)  acids,  and  (3)  bases. 

1.  Preparation  from  water.  Hydrogen  may  be  liberated  from  water 
by  two  general  methods :  (a)  by  the  action  of  the  electric  current,  as 
described  under  the  methods  for  the  preparation  of  oxygen  (see 
p.  17) ;  (b)  by  the  action  of  certain  metals.  A  few  metals,  such 
as  sodium,  potassium,  and  calcium,  act  rapidly  upon  water  even  at 
ordinary  temperatures,  liberating  one  half  of  the  hydrogen  present 
in  the  water.  The  remainder  of  the  hydrogen,  together  with  all  the 

35 


36 


GENERAL  CHEMISTRY 


•1 


C=o 


oxygen,  combines  with  the  metal  to  form  compounds  which  belong  to 
the  class  known  as  bases,  or  hydroxides.  The  reaction  which  takes  place 
when  water  is  decomposed  by  sodium  may  be  represented  as  follows : 

sodium  -f  water >•  sodium  hydroxide  +  hydrogen 

[hydrogen"]  podium 

oxygen     J  hydrogei 

[oxygen 

Fig.  15  represents  a  form  of  apparatus  used  in  preparing  hydrogen  by  this 
method.  A  pellet  of  sodium  is  pushed  into  the  end  of  a  short  piece  of  lead  or  tin 
pipe,  the  other  end  of  the  pipe  being  hammered  until 
closed.  The  pipe  containing  the  sodium  is  then 
dropped  into  a  trough  of  water  arranged  as  shown 
in  the  figure.  The  hydrogen  liberated  by  the  action 
of  the  sodium  upon  the  water  rises  in  bubbles  and  is 
caught  in  the  jar.  The  sodium  hydroxide  formed  in 
the  reaction  may  be  recovered  by  evaporating  the 
water  remaining  after  the  action  is  completed. 

Other  metals,  such  as  magne- 
sium and  iron,  also  decompose 
water  rapidly  but  only  at  high 
temperatures.  In  such  cases  the 
decomposition  is  best  effected  by 
heating  the  metal  to  redness  in.  a 
glass  or  porcelain  tube  and  then  passing  steam  over  the  red-hot  metal. 
Under  these  conditions  the  metal  combines  with  the  oxygen  of  the 
water  to  form  an  oxide,  while  the  hydrogen  is  liberated.  For  example, 
in  the  case  of  iron  the  reaction  may  be  represented  as  follows : 

iron  +  water >•  iron  oxide  +  hydrogen 

["hydrogen"]         Tiron      ~| 
[oxygen     J         [oxygenj 

Fig.  16  represents  a  simple  form  of  apparatus  which  may  be  used  in  prepar- 
ing hydrogen  by  this  method.  Iron  in  the  form  of  fine  wire  or  tacks  is  placed  in 
the  tube  A  and  heated  to  a  high  temperature.  The  water  in  flask  B  is  then 
boiled  and  the  resulting  steam  passed  into  A.  The  iron  in  the  tube  combines 
with  the  oxygen  of  the  steam  to  form  oxide  of  iron,  which,  being  a  solid  sub- 
stance, remains  in  the  tube,  while  the  hydrogen  passes  on  and  is  collected  in 
the  receiver  C,  as  is  shown  in  the  diagram. 

2.  Preparation  from  acids.  A  more  convenient  method  for  preparing 
hydrogen  in  the  laboratory  consists  in  liberating  it  from  acids  by  the 
action  of  metals.  For  this  purpose  any  of  the  metals  which  liberate 
hydrogen  from  water,-  but  only  these,  may  be  employed.  Usually  zinc 


FIG. 15 


HYDKOGEN 


37 


or  iron  is  used.  The  acids  commonly  employed  are  either  hydro- 
chloric acid  or  sulfuric  acid.  The  former  is  an  aqueous  solution  of 
a  gaseous  compound  containing  2.76  per  cent  of  hydrogen  and  97.24 
per  cent  of  chlorine,  while  the  latter  is  an  aqueous  solution  of  an  oily 


FIG. 16 

liquid  containing  2.06  per  cent  of  hydrogen,  32.69  per  cent  of  sulfur, 
and  65.25  per  cent  of  oxygen.  To  liberate  hydrogen  it  is  only  neces- 
sary to  bring  the  acid,  properly  diluted  with  water,  into  contact  with 
the  metal.  The  metal  gradually  passes  into  solution,  while  the  hydro- 
gen of  the  acid  is  in  turn  set  free.  The  liberation  of  the  hydrogen  is 
indicated  by  the  effervescence  of  the  liquid.  When  zinc  and  sulfuric 
acid  are  used  in  the  preparation,  the  reaction  may  be  represented  in  a 
general  way  as  follows : 


zinc  H-  sulfuric  acid 

[hydrogen! 
sulfur 
oxygen 


zinc  sulf ate  +  hydrogen 

pine      "I 

sulfur 
[oxygenj 


It  will  be  noted  that  the  zinc  simply  takes  the  place  of  the  hydro- 
gen in  the  acid.  The  resulting  compound  of  zinc,  sulfur,  and  oxygen, 
known  as  zinc  sulfate,  is  a  white  solid  which  remains  dissolved  in 
the  water  present  and  may  be  obtained  by  evaporating  the  solution. 

When  iron  and  hydrochloric  acid  are  used  in  the  preparation  of 
hydrogen,  the  reaction  may  be  represented  as  follows : 

iron  -|-  hydrochloric  acid >•  iron  chloride  +  hydrogen 

["hydrogen"!  Rron        "1 

^chlorine  J  LchlorineJ 


38 


GENERAL  CHEMISTRY 


FIG. 17 


A  convenient  form  of  apparatus  for  preparing  hydrogen  by  the  action  of  metals 
upon  acids  is  shown  in  Fig.  17.  The  metal  is  placed  in  flask  A,  which  is  fitted 
with  a  cork  and  connected  with  tubes,  as  shown  in  the 
figure.  The  acid,  properly  diluted  with  water,  is  added  a  little 
at  a  time  through  the  funnel  tube  B.  The  liberated  hydrogen 
escapes  through  C  and  is  collected  in  receivers,  as  shown  in 

the  figure.  The  hydrogen 
which  first  escapes  through 
the  exit  tube  is  mixed  with 
the  air  originally  present  in 
flask  A.  Such  a  mixture  of 
hydrogen  and  air  is  violently 
explosive  when  brought  in 
contact  with  aflame.  There- 
fore one  must  keep  all 
flames  away  from  the  ap- 
paratus. Moreover,  one  should  not  collect  the  hydrogen  until  a  sufficient  amount 
of  it  is  generated  to  displace  all  the  air  previously  contained  in  the  flask. 

A  more  convenient  form  of  apparatus  to  use  is  that  shown  in  Fig.  18.  It  is 
known  as  a  Kipp  generator  and  has  the  advantage  of  being  automatic  in  its 
action.  The  metal  is  placed  in  A ,  and  the  acid  poured  into 
B.   When  the  stopcock  D  is  opened,  the  acid  runs  down 
into  C  and  up  into  A,  where  it  comes  in  contact  with 
the  metal.    The  hydrogen  generated  escapes  through  D. 
If  now  the  stopcock  is  closed,  the  hydrogen,  being  unable 
to  escape  through  the  tube,  pushes  the  acid  away  from 
the  metal  in  A  down  into  C  and  up  into  B,  so  that  the 
action  ceases  until  the  stopcock  is  again  opened. 

3.  Preparation  from  bases.  The  bases,  or  metallic 
hydroxides,  are  compounds  of  a  metal  with  oxy- 
gen and  hydrogen.  Certain  metals,  such  as  zinc 
and  aluminium,  react  with  some  of  the  bases, 
liberating  the  hydrogen  present  in  them.  Thus, 
when  zinc  is  added  to  a  solution  of  sodium  hydrox- 
ide and  the  mixture  is  heated,  the  zinc  takes  the  place  of  the  hydro- 
gen of  the  base.  The  reaction  may  be  represented  as  follows: 


FIG.  18 


zinc  H-  sodium  hydroxide 

fsodium    ~j 

oxygen 
[hydrogenj 


sodium  zincate  -h  hydrogen 

("sodium" 
I  oxygen 
[zinc 


Commercial  method  of  preparation.  Hydrogen  is  not  prepared  to 
any  great  extent  on  a  commercial  scale,  since  the  cheaper  coal  gas 
or  natural  gas  serves  most  purposes  equally  well.  When  wanted  in 


HYDROGEN  39 

large  quantities  it  is  prepared  by  the  action  of  sulfuric  acid  upon  iron, 
or  by  the  decomposition  of  water  by  the  electric  current.  The  latter 
method  is  economical  only  when,  cheap  water  power  is  available  for 
generating  the  electric  current. 

Properties.  Hydrogen,  like  oxygen,  is  a  colorless,  odorless,  and 
tasteless  gas.  It  has  the  greatest  specific  heat,  as  well  as  the  highest 
thermal  conductivity,  of  any  gas.  One  liter  of  it  weighs  0.08987  g. 
It  is  the  lightest  of  all  known  substances,  being  14.385  times  lighter 
than  ah- ;  it  may  therefore  be  transferred  from  one  vessel  to  another 
by  pouring  it  upward,  as  shown  in  Fig.  19.  The  hydrogen  in  the 
cylinder  A  rises  to  the  top  of  the  cylinder  B  and  forces  the  air  out. 
The  solubility  of  hydrogen  in  water  is  very 
small,  being  only  about  one  half  as  great  as 
that  of  oxygen. 

Dewar  was  the  first  to  obtain  hydrogen  in  the 
liquid  state.  He  cooled  the  gas  to  a  temperature 
of  —205°  by  means  of  liquid  air,  and  at  the 
same  time  subjected  it  to  a  pressure  of  180  at- 
mospheres. In  this  way  it  was  obtained  as  a 
colorless,  transparent  liquid,  boiling  at  —252.7°  FIG.  19 

under   a  pressure    of   1    atmosphere.     This   is 

the  lightest  liquid  known,  having  a  density  of  but  0.07  at  its  boiling 
point.  When  liquid  hydrogen  is  evaporated  under  very  small  pressure, 
solid  hydrogen  is  obtained  as  a  transparent,  snowlike  body  melting 
at  about  —  259°. 

A  large  number  of  metals  have  the  property  of  absorbing  or "  occlud- 
ing" hydrogen.  The  quantity  so  absorbed  by  most  of  the  metals  is  not 
large,  but  a  few,  such  as  gold,  platinum,  and  especially  palladium 
take  up  large  volumes  of  the  gas.  The  quantity  absorbed  varies  not 
only  with  the  metal  but  also  with  the  physical  condition  of  the  metal, 
as  well  as  with  the  temperature  and  pressure  under  which  the  ab- 
sorption takes  place.  One  volume  of  palladium,  in  the  form  of  a  pow- 
der, at  ordinary  temperatures  absorbs  over  800  volumes  of  the  gas. 
It  is  because  of  this  property  that  hydrogen,  when  conducted  into 
hot  tubes  made  of  iron  or  platinum,  passes  through  the  walls  of  the 
tube  to  a  considerable  extent. 

The  absorption  of  hydrogen  by  palladium  can  be  strikingly  shown  by  using 
strips  of  this  metal  as  electrodes  in  the  decomposition  of  water  by  the  electric 
current.  When  the  circuit  is  closed,  oxygen  is  at  once  evolved  at  one  electrode 


40  GENERAL  CHEMISTRY 

and  hydrogen  at  the  other.  It  will  be  noted,  however,  that  the  amount  of  hydro- 
gen evolved  at  first  is  relatively  small,  but  gradually  increases  as  the  palladium 
becomes  saturated  with  the  gas. 

Chemical  conduct.  Although  hydrogen  is  quite  inactive  at  ordinary 
temperatures,  nevertheless  under  proper  conditions  it  combines  directly 
with  many  of  the  elements,  and  even  decomposes  some  compounds  by 
uniting  with  certain  elements  present  in  them.  Just  as  the  compounds 
of  oxygen  with  any  other  one  element  are  termed  oxides,  so  those 
containing  hydrogen  in  combination  with  another  element  are,  as  a 
class,  known  as  hydrides.  Many  of  the  individual  members  of  this 
class  of  compounds,  however,  have  other  names  applied  to  them. 

1.  Action  of  hydrogen  upon  elements.  At  suitable  temperatures  hydro- 
gen combines  directly  with  nitrogen  to  form  the  gaseous  compound 
known  as  ammonia,  with  sulfur  to  form  the  foul-smelling  gas  known 
as  hydrogen  sulfide,  with  chlorine  to  form  hydrogen  chloride,  a  gas 
whose  solution  in  water  is  termed  hydrochloric  acid.  The  union  with 
chlorine  can  also  be  brought  about  by  strong  light,  a  mixture  of  hy- 
drogen and  chlorine  exploding  with  great  violence  when  exposed  to 
direct  sunlight.  Hydrogen  also  combines  with  a  number  of  the  metals. 
It  is  characterized  especially,  however,  by  its  affinity  for  oxygen,  with 
which  it  combines  to  form  water.  Experiments  show  that  the  ratio  in 
which  these  two  gases  combine  is  1  of  hydrogen  to  7.94  of  oxygen  by 
weight,  or  2.0024  of  hydrogen  to  1  of  oxygen  by  volume.  A  large 
amount  of  heat  is  set  free  in  this  reaction,  amounting  to  34,215  cal. 
for  each  gram  of  hydrogen  entering  into  combination. 

The  union  of  hydrogen  and  oxygen  probably  takes  place  at  ordinary  tempera- 
tures, but  the  speed  of  the  reaction  is  so  slow  that  no  change  can  be  detected 
even  after  long  intervals  of  time  (see  p.  27).  As  the  temperature  is  raised  the  speed 
increases.  Thus  Meyer  and  Raum  found  that  the  two  gases,  when  mixed  in  the 
proportion  of  two  volumes  of  hydrogen  to  one  volume  of  oxygen  and  heated  to 
100°  for  218  days  showed  no  appreciable  combination.  When  heated  to  300°  for 
65  days  it  was  found  that,  in  different  trials,  from  0.4  per  cent  to  9.5  per  cent  of 
the  mixture  had  combined.  At  500°  the  change  is  still  more  marked,  but  takes 
place  gradually  and  requires  several  hours  for  completion.  'At  a  temperature 
roughly  approximating  800°  the  union  of  the  two  takes  place  with  explosive 
violence.  The  temperature  at  which  this  instantaneous  combination  takes  place 
is  constant  when  the  conditions  are  exactly  the  same.  It  is  modified,  however,  by 
very  slight  changes  in  these  conditions,  due  to  the  catalytic  effect  of  foreign  sub- 
stances, such  as  moisture  and  the  materials  of  which  the  tube  containing  the 
gases  is  made.  Certain  catalyzers,  such  as  finely  divided  platinum,  bring  about 
practically  instantaneous  combination  at  ordinary  temperatures. 


HYDEOGEN 


41 


The  union  of  hydrogen  and  oxygen,  and  the  resulting  formation 
of  water,  is  best  shown  by  burning  hydrogen  in  oxygen  or  air.  A 
convenient  apparatus  is  shown  in  Fig.  20. 

The  hydrogen  is  generated  in  flask  A  and  passed  through  the  tube  B  filled  witH 
porous  calcium  chloride,  which  removes  the  moisture  from  the  gas.  After  the  air 
has  been  displaced  from  the  apparatus,  the  hydrogen  escaping  at  the  jet  Cis  ignited. 
Almost  instantly  a  dewlike  substance  is  deposited  on  the  cold  sides  of  the  jar  D. 
This  may  be  collected  and  proved  to  be  water. 


FIG.  20 


The  hydrogen,  if  pure,  burns  with  a  colorless  but  very  hot  flame. 
When  burned  in  air  much  of  the  heat  resulting  from  the  union  of  the 
hydrogen  and  oxygen  is  absorbed  in  heating  the  inert  nitrogen  present. 
To  obtain  the  maximum  temperature  one  must  use  pure  oxygen 
instead  of  air.  Moreover,  the 
hydrogen  and  oxygen  must  be 
brought  together  in  the  exact 
proportion  in  which  they  unite ; 
otherwise  the  gas  which  is  left 
uncombined  will  absorb  a  por- 
tion of  the  heat.  An  apparatus 
arranged  for  burning  hydrogen 
in  this  way  is  known  as  the  oxy- 
hydrogen  blowpipe.  This  con- 
sists primarily  of  two  tubes 
(Fig.  21),  one  inside  the  other.  Hydrogen  is  forced  in  through  the 
tube  H  and  ignited  at  the  end  of  tube  A.  Oxygen  is  then  forced  -in 
through  the  tube  0.  The  two  gases  are  thus  brought  together  at  the 
tip  of  the  tube  A,  and  combine  with  the  evolution  of  a  large  amount  of 


FIG.  21 


42 


GENERAL  CHEMISTRY 


FIG.  22 


heat.  Under  most  favorable  conditions  a  temperature  of  about  2500° 
may  thus  be  obtained.  Ordinarily,  however,  the  temperature  reached 
does  not  exceed  1800°.  At  the  high  temperature  of  the  flame  a  piece 
of  lime  glows  intensely,  while  an  iron  wire  burns  with  the  greatest 
brilliancy.  Platinum  (melting  point  1755°)  is  readily  melted  when 
exposed  to  it. 

While  it  is  thus  possible  to  burn "  hydrogen 
safely  in  air  or  pure  oxygen  by  limiting  the 
amounts  of  the  two  gases  brought  in  contact 
with  each  other  at  any  instant,  mixtures  of  the 
gases  in  any  considerable  quantities  explode 
with  terrific  violence  when  ignited. 

_  This  fact  may  be  demonstrated  without  danger  in 
the  following  way:  A  bell  jar  of  two  or  three  liters 
capacity  is  fitted  with  a  cork  provided  with  a  short  open 
tube  of  about  1  cm.  diameter,  as  shown  in  Fig.  22.  The  tube  is  closed  with  a  small 
rubber  stopper,  and  the  bell  jar  filled  with  hydrogen,  the  gas  being  collected  over 
water.  When  entirely  filled  the  jar  is  set  on  blocks  of  wood,  the  stopper  removed, 
and  the  hydrogen  ignited  at  the  top  of  the  tube.  As  the  hydrogen  rises  through 
the  tube  and  is  burned,  air  enters  the  jar  from  below  and  mixes  with  the  remain- 
ing hydrogen.  When  a  volume  sufficient  to  form  an  explosive  mixture  with 
the  hydrogen  has  thus  entered  the  jar,  a  violent  explosion  results.  Since  the 
jar  is  open  at  the  bottom,  and  the  intensity  of  the  explosion  is  diminished 
by  the  presence  of  the  nitrogen  in  the  air,  there  is  no  danger  attending  the 
experiment.  The  ex- 
plosion of  the  two 
gases  reaches  the 
maximum  effect  with 
mixtures  of  pure  hy- 
drogen and  oxygen  in 
the  exact  proportion 
in  which  they  com- 
bine. Such  mixtures 
should  never  be  ex- 
ploded except  in  small 
quantities  and  by  ex- 
perienced chemists. 


FIG.  23 


2.  Action  of  hydrogen  upon  compounds.  Hydrogen  not  only  combines 
directly  with  many  elements,  such  as  chlorine  and  oxygen,  when  pres- 
ent in  the  free  state,  but  under  favorable  conditions  it  will  remove 
these  elements  from  some  of  their  compounds. 

This  action  may  be  shown  by  introducing  some  black  oxide  of  copper  into  a 
hard  glass  tube  C  (Fig.  23),  which  is  connected  with  a  hydrogen  generator  A 


HYDKOGEN  43 

and  drier  B,  as  shown  in  the  figure.  Hydrogen  is  generated,  and,  after  the  air  has 
been  completely  displaced  from  the  entire  apparatus,  the  tube  containing  the  oxide  of 
copper  is  heated.  The  hydrogen  combines  with  the  oxygen  present  in  the  copper 
oxide  to  form  water,  which  condenses  in  the  colder  portions  of  the  tube  near  the 
end,  while  the  black  color  of  the  oxide  of  copper  gradually  gives  way  to  the  red- 
dish tint  of  copper  as  the  action  progresses.  The  change  may  be  represented  as 

follows : 

hydrogen  +  copper  oxide >-  water  +  copper 

"copper  "j          ("hydrogen"! 
.  oxygen J          [oxygen     J 

Many  other  metallic  oxides,  such  as  that  of  iron,  may  be  substituted  for  the 
oxide  of  copper  in  this  experiment.  When  oxide  of  iron  is  used,  the  change  may 
be  represented  as  follows  : 


hydrogen  +  iron  oxide >-  water  +  iron 

Tiron      ~\        Phydrogenl 
LoxygenJ        [_°xygei1     J 


Reduction;  reducing  agent.  When  oxygen  is  removed  from  a  com- 
pound, the  change  is  known  as  reduction.  The  compound  from  which 
the  oxygen  is  removed  is  said  to  be  reduced,  while  the  substance  which 
unites  with  the  oxygen  is  called  the  reducing  agent.  Thus,  in  the  ex- 
periment with  the  oxide  of  copper  the  hydrogen  is  termed  the  reduc- 
ing agent  and  the  copper  oxide  is  said  to  be  reduced.  It  will  be  observed 
that  reduction  is  just  the  reverse  of  oxidation.  In  the  latter  process 
oxygen  is  added  to  a  substance,  and  in  the  former  it  is  taken  away. 

The  reaction  between  hydrogen  and  oxide  of  iron.  It  will  be  observed 
that  the  reaction  which  takes  place  when  hydrogen  is  passed  over  heated 
iron  oxide  (see  above)  is  just  the  reverse  of  that  which  takes  place 
when  steam  is  passed  over  red-hot  iron  (see  p.  36).  In  the  one 
case  hydrogen  and  iron  oxide  react  to  form  water  and  iron ;  in  the 
other  the  iron  and  water  react  to  form  iron  oxide  and  hydrogen.  The 
two  reactions  may  be  expressed  as  follows : 

iron  4-  water  (steam)  <     >  iron  oxide  4-  hydrogen 

This  reaction  is  similar  to  that  which  takes  place  in  the  conversion  of 
oxygen  into  ozone,  and  of  barium  oxide  into  barium  peroxide,  in  that 
it  is  reversible.  Under  the  same  set  of  conditions  the  two  opposing 
reactions  may  go  on  until  an  equilibrium  is  reached.  It  is  possible, 
however,  to  devise  conditions  under  which  one  of  the  reactions  is  so 
much  aided  as  to  practically  reach  completion.  This  subject  will  be 
more  sytematically  discussed  in  Chapter  XVI. 


44  GENERAL  CHEMISTRY 

Action  of  hydrogen  upon  the  system  when  inhaled.  Pure  hydrogen 
is  not  poisonous  and  may  be  breathed  without  danger,  but  the  sulfuric 
acid  and  zinc  used  in  its  preparation  frequently  contain  small  amounts 
of  arsenic,  and  the  hydrogen  generated  from  such  substances  is  mixed 
with  a  gaseous  compound  of  arsenic  and  hydrogen,  which  is  exceed- 
ingly poisonous  and  must  not  be  inhaled. 

Uses  of  hydrogen.  Hydrogen  is  used  commercially  only  to  a  limited 
extent.  Formerly  considerable  quantities  of  it  were  used  in  the  oxy- 
hydrogen  blowpipe.  The  electric  current,  however,  has  proved  to  be 
a  cheaper  as  well  as  a  more  convenient  source  of  heat.  It  is  also 
employed  as  a  reducing  agent,  especially  in  the  process  of  refining 
certain  oils.  Because  of  its  extreme  lightness  it  is  used  for  inflating 
dirigible  airships.  The  cheaper  coal  gas,  although  heavier  than 
hydrogen,  is  used  for  inflating  ordinary  balloons. 


CHAPTER  IV 


PROPERTIES  OF  GASES 

The  gaseous  state  a  property  of  all  substances.  Experiment  has 
shown  that  every  substance  tends  to  pass  into  the  gaseous  state  if 
its  temperature  is  raised  sufficiently,  but  it  is  not  always  possible  to 
actually  bring  about  such  a  change.  In  some  cases  the  temperature 
required  is  so  high  that  it  cannot  be  attained  by  laboratory  methods ; 
in  others  the  substance  decomposes  before  the  required  temperature 
is  reached.  For  example,  it  will  be  recalled  that  mercuric  oxide  and 
potassium  chlorate  yield  oxygen  when  heated,  while  sugar  decomposes 
into  carbon,  water  vapor,  and  other  products. 

Experience  also  shows  that  at  some  lower  temperature  every  gas 
condenses  to  a  liquid  or  a  solid.  Fifty  years  ago,  when  the  laws  gov- 
erning gases  were  not  so  well  known  as  they  are  now  (see  p.  76), 
some  gases,  including  oxygen,  nitrogen,  and  hydrogen,  resisted  all 
efforts  directed  toward  their  liquefaction,  and  these  were  termed  per- 
manent gases.  All  known  gases  have  now  been  liquefied,  and  the 
gaseous  state  is  recognized  as  a  general  condition  into  which  all  matter 
may  pass,  and  not  a  peculiarity  of  any  particular  group  of  substances. 

Characteristics  of  the  gaseous  state.  1.  Expansibility.  The  most 
striking  characteristic  of  a  gas  is  its  tendency  to  expand  indefinitely, 
so  as  to  distribute  itself  uniformly  through- 
out all  the  space  in  which  it  is  confined. 
If  the  gas  is  suddenly  set  free  in  a  vacuum, 
this  distribution  occurs  with  great  rapidity. 

This  rapid  distribution  of  a  gas  throughout  a 
vacuum  may  be  very  successfully  demonstrated  by 
the  following  experiment :  A  small,  thin-walled 
bulb  A  (Fig.  24)  is  filled  with  bromine,  sealed, 
and  placed  in  a  glass  vessel  B,  which  can  be  closed 
and  exhausted  by  a  water  pump  attached  at  C. 
A  glass  rod  is  arranged  to  slip  through  a  tight- 
fitting  rubber  stopper  D.  When  the  vessel  has  been  exhausted,  the  rod  is  pushed 
down  so  as  to  crush  the  bulb.  Almost  instantly  the  reddish  vapor  of  bromine 
can  be  seen  in  every  part  of  the  vessel. 

45 


.Fie.  24 


46  GENERAL  CHEMISTRY 

If  the  gas  is  set  free  in  a  space  already  occupied  by  another  gas, 
the  distribution  takes  place  very  slowly.  The  one  gas  is  said  to  diffuse 
through  the  other,  which  presents  a  certain  obstacle  to  its  spread. 
Sooner  or  later,  however,  the  two  become  evenly  mixed.  The  vapor 
resulting  from  the  evaporation  of  a  few  drops  of  bromine  may  gradually 
be  perceived  by  its  odor  in  every  part  of  a  large  room. 

Diffusion  will  take  place  quite  independently  of  the  relative  weights 
of  the  two  gases.  We  may  introduce  a  heavy  gas  under  a  lighter  one 
in  a  closed  vessel,  but  diffusion  will  take  place,  notwithstanding  the 
difference  in  density.  Likewise  two  gases  of  unequal  density,  when 
once  mixed,  show  no  tendency  to  separate  into  layers,  with  the  heavier 
one  below  the  lighter.  Complete  and  permanent  diffusion  is  charac- 
teristic of  all  gas  mixtures. 

2.  Compressibility.  A  second  and  equally  important  characteristic 
of  gases  is  their  compressibility.  Liquids  and  solids  are  very  little 
affected  by  pressure,  but  the  volume  of  a  gas  is  very  greatly  changed 
by  comparatively  small  changes  in  pressure.  A  very  familiar  illus- 
tration of  this  is  found  in  the  common  experience  of  pumping  up  a 
bicycle  or  automobile  tire.  A  surprising  volume  of  air  may  be 
pumped  into  what  is  an  almost  unchanged  volume  in  the  tire. 

Obviously,  before  we  can  get  a  clear  idea  as  to  the  density  of  a  gas, 
that  is,  the  mass  of  unit  volume,  or  the  concentration,  we  must  have 
an  understanding  of  the  way  in  which  the  volume  of  a  given  mass 
changes  with  the  various  physical  conditions,  and  we  must  adopt  a  set 
of  standard  conditions  under  which  all  measurements  are  to  be  made. 

The  gas  laws.  A  large  number  of  laws  relating  to  gases  have  been 
formulated,  and  many  of  these  will  be  mentioned  in  subsequent 
pages.  At  present  it  will  be  sufficient  to  consider  four,  which  will 
serve  to  define  the  conduct  of  gases  under  the  usual  variety  of 
conditions. 

1.  The  relation  of  volume  to  pressure;  the  law  of  Boyle.  In  1660 
Robert  Boyle,  an  Englishman  of  remarkable  scientific  accuracy  for  his 
time,  made  some  careful  measurements  upon  the  compressibility  of 
gases.  He  succeeded  in  establishing  the  generalization  known  as 
Boyle's  law,  which  may  be  stated  thus :  The  volume  which  a  given 
mass  of  a  gaseous  substance  occupies  is  inversely  proportional  to  the 
pressure  under  which  it  is  measured,  provided  the  temperature  remains 
constant.  Doubling  the  pressure  diminishes  the  volume  one  half; 
diminishing  the  pressure  one  half  doubles  the  volume.  The  product 


PEOPERTIES  OF  GASES  47 

of  the  pressure  into   the  volume  is  therefore   constant.  Stated  in 
algebraic  form,  we  have  the  equation 

-C  (1) 


in  which  PI  and  V^  are  the  pressure  and  volume  of  a  given  mass  of 
gas  under  one  set  of  conditions,  and  P2  and  V2  those  under  another, 
C  being  their  constant  product.  The  magnitude  of  C  will  of  course 
depend  upon  the  quantity  of  gas  taken  for  experimentation. 

Like  most  of  the  laws  of  science,  Boyle's  law  is  only  an  approximate 
statement  of  the  facts,  since  all  gases  do  not  act  in  precisely  the  same 
way,  and  extreme  conditions  introduce  irregularities.  For  example, 
hydrogen  gas  is  not  as  compressible  as  the  law  leads  us  to  expect; 
at  very  high  pressures  all  gases  resist  pressure  more  than  at  moderate 
pressures ;  when  increase  of  pressure  brings  the  gas  near  its  point  of 
liquefaction,  it  is  more  easily  compressed  than  the  law  predicts.  Within 
the  range  of  accuracy  required  for  most  chemical  purposes,  however, 
Boyle's  law  is  a  remarkably  exact  statement  of  the  facts  and  holds 
true  for  all  gases. 

Standard  pressure.  For  practical  purposes  we  must  choose  some 
standard  pressure  to  which  all  gas  volumes  are  to  be  referred.    This 
P  is  most  conveniently  chosen  as  the  average  pressure  of  the 

atmosphere  at  the  sea  level.  This  is  equal  to  the  pressure 
exerted  by  a  column  of  mercury  760  mm.  in  height,  or  to 
1033  g.  per  square  centimeter. 

The  graphic  representation  of  Boyle's  law.  Relations  like  those 
stated  in  Boyle's  law  are  best  represented  by  a  curved  line  every 
point  on  which  is  determined  by  its  distance 
from  two  standard  lines  at  right  angles  to 
each  other.  These  are  called  the  axes,  or 
coordinates,  the  horizontal  one  being  named 

t^  ~tT          ij  ""  v      the  abscissa  and  the  vertical  one  the  ordinate. 

-p      2e:  If  we  measure  volume  along  the  abscissa  and 

pressure  along  the  ordinate,  we  can  draw  a 

curve  like  that  represented  in  Fig.  25.  The  areas  p^v^  p2v2,  p3v3,  are  equal,  and 
the  curve  AB  is  known  in  geometry  as  a  rectangular  hyperbola.  The  distance 
of  the  curve  from  the  axes  is  evidently  determined  by  the  numerical  value  of  the 
product  C  in  equation  (1),  that  is,  by  the  area  p^yv  p2v2,  etc.  This  is  in  turn 
dependent  on  the  quantity  of  gas  under  observation. 

2.  The  relation  of  volume  to  temperature.  The  law  connecting  volume 
with  temperature  was  formulated  independently  in  1801  by  the  French 
chemist  Gay-Lussac  and  the  English  schoolmaster  Dalton.  It  is  also 


48  GENEEAL  CHEMISTEY 

sometimes  called  the  law  of  Charles.  These  investigators  found  that  all 
gases  expand  when  the  temperature  is  raised  (the  pressure  being  held 
constant),  and  that  equal  volumes  of  all  gases  expand  to  the  same 
extent  for  a  given  increase  in  temperature.  Let  us  suppose  that  the 
volume  of  the  gas  has  been  measured  at  zero  on  the  centigrade  scale. 
Experiment  shows  that  a  rise  of  one  degree  causes  an  expansion  equal 
to  ^3  of  this  volume ;  a  rise  of  two  degrees,  an  expansion  of  ^|^. 
That  is  to  say,  if  273  cc.  of  the  gas  is  measured  at  zero,  the  volume 
at  1°  above  will  be  274  cc. ;  at  2°  above,  275  cc.  At  1°  below  it  will 
contract  to  272  cc.,  and  at  2°  below,  to  271  cc.  If  the  same  rate  of 
contraction  holds  good  at  all  temperatures,  then  at  —  272°  the  volume 
will  be  1  cc.,  and  at  —  273°  the  volume  will  be  zero.  Obviously  this 
last  conclusion  cannot  be  true,  but  it  must  mean  that  before  such  a 
temperature  is  reached,  all  gases  will  have  become  liquids,  in  which 
state  the  law  will  not  apply.  This  interpretation  is  borne  out  by  the 
fact  that  helium,  the  most  difficult  of  all  gases  to  liquefy,  passes  into 
a  liquid  at  -  268.7°. 

The  absolute  scale  of  temperature.  If  we  were  to  construct  a  ther- 
mometer having  divisions  of  the  same  size  as  those  on  the  centigrade 
scale,  but  with  the  zero  point  at  —  273°  on  the  latter  scale,  then  the 
point  at  which  water  freezes  would  be  273°.  At  272°  on  this  scale 
the  273  cc.  of  gas  discussed  in  the  last  paragraph  would  measure 
272  cc. ;  at  271°,  271  cc. ;  at  1°,  1  cc.  On  such  a  scale  the  volume 
of  the  gas  would  be  proportional  to  the  temperature  at  every  point. 
This  is  known  as  the  scale  of  absolute  temperatures,  the  point  —  273° 
on  the  centigrade  scale  being  absolute  zero.  Evidently  the  absolute  tem- 
perature may  be  obtained  by  adding  273°  to  the  centigrade  reading. 

The  law  of  Gay-Lussac  (or  of  Charles).  The  law  established  by 
Gay-Lussac  may  now  be  stated  in  the  following  quantitative  form: 
The  volume  occupied  by  a  given  mass  of  a  gas  at  different  tempera- 
tures is  proportional  to  the  absolute  temperature  (pressure  remaining 
constant).  If  Vl  and  F2  are  the  volumes  at  the  absolute  temperatures 
T{  and  T2,  then  y  T 

This  equation  may  be  combined  with  equation  (1),  which  expresses 
Boyle's  law,  giving  the  equation 

(3) 


PROPERTIES  OF  GASES  49 

This  expresses  in  one  equation  the  variation  in  volume  due  to  both 
pressure  and  temperature. 

Standard  conditions.  To  the  standard  pressure  already  adopted  we 
must  now  add  a  standard  temperature,  and  this  is  chosen  as  0°  centi- 
grade (equal  to  273°  absolute).  If  we  designate  the  volume  under 
standard  conditions  by  V#  then  V#  760,  and  273  will  be  the  values  of 
volume,  pressure,  and  temperature  under  one  standard  set  of  conditions 
and  will  correspond  to  one  set  of  values  in  equation  (3).  Substitut- 
ing them  in  place  of  V^  P^  T^  and  dropping  the  subscript  of  the 
values  F2,  P^  T2  as  no  longer  necessary,  we  get  the  equation 

PXFX273 
760  x  T 

in  which  P,  F,  and  T  are  the  values  under  any  set  of  conditions  other 
than  those  adopted  as  standard.  For  example,  if  in  preparing  oxygen*  a 
volume  of  685  cc.  happened  to  be  obtained  at  a  temperature  of  22°  and 
under  a  pressure  of  750  mm.,  the  volume  under  standard  conditions 

wouldbe  750  x  685  x  273 

s  =  760  x  (273  +  22)  = 

If  we  employ  the  centigrade  scale  we  shall  have  to  state  the  law  of  Gay-Lussac 
in  the  following  way :  A  gas  measured  at  0°  changes  its  volume  by  the  fraction 
5|T  for  every  degree  that  the  temperature  varies  from  zero.  Designating  the 
volume  at  zero  by  V0,  and  at  a  different  temperature  by  Vt, 

Fo(l  +  0.003660 

If  this  equation  is  combined  with  (1),  which  states  Boyle's  law,  the  following 
equation  is  obtained  : 

y   - "  x  Vt  x-v 

760(1+  0.00366  t) 

The  values  used  in  the  example  in  the  last  paragraph,  when  substituted  in  this 
equation,  will  give  the  same  result  as  with  equation  (4). 

3.  The  relation  of  temperature  to  pressure.  We  have  seen  that  when 
the  pressure  is  maintained  constant,  the  volume  increases  in  propor- 
tion to  the  absolute  temperature.  From  this,  together  with  Boyle's 
law,  it  follows  that  if  the  volume  is  kept  constant,  the  pressure  will 
increase  in  proportion  to  the  absolute  temperature.  Gay-Lussac  made 
some  experiments  upon  this  point  and  found  that  the  pressure  does 
so  increase.  This  principle  is  not  of  as  frequent,  application  in  chemi- 
cal calculations  as  the  more  familiar  law  of  Gay-Lussac,  but  in  many 


50 


GENERAL  CHEMISTRY 


lines  of  mechanics  it  is  of  importance.  For  example,  it  enables  us  to 
calculate  the  pressure  in  a  steam  boiler  at  a  temperature  £,  if  we  know 
the  value  at  some  other  temperature  tr.  By  measurement  of  the  pres- 
sure reached  in  an  explosion  we  may  also  calculate  the  temperature, 
or  vice  versa. 

4.  The  rate  of  diffusion;  the  law  of  Graham.  In  1883  the  Scottish 
chemist  Thomas  Graham  made  a  series  of  studies  upon  what  he  called 
the  rate  of  diffusion  of  gases,  by  which  he  meant  the  rate  at  which 
various  gases  will  pass  through  a  minute  pinhole,  or  through  porous 
materials  such  as  unglazed  pottery.  As  a  result  of  his  experiments 
he  found  that  under  definite  conditions  of  temperature  and  pressure 
the  rate  of  diffusion  is  inversely  proportional  to  the  square  root  of  the 

density  of  the  gas.  Of  two  gases  the 
lighter  will  therefore  diffuse  the  more 
rapidly.  Oxygen  and  hydrogen  have 
densities  which  are  almost  exactly  in 
the  ratio  16  : 1,  and  their  rates  of  diffu- 
sion are  therefore  in  the  ratio 


FIG.  26 


In  other  words,  hydrogen  will  leak 
through  fine  pores  four  times  as  fast 
as  oxygen. 

Demonstration  of  diffusion.  This  property 
may  be  demonstrated  by  the  use  of  the 
apparatus  represented  in  Fig.  26.  A  small 
battery  jar  A  is  connected  with  a  tight- 
fitting  rubber  stopper  or  plaster-of-Paris  joint 
with  a  glass  tube  B,  the  other  end  of  which  passes  just  through  a  stopper  in  the 
vessel  C.  The  vessel  is  half  filled  with  water  and  is  provided  with  a  second 
tube  D,  drawn  to  a  small  jet  at  E  and  extending  to  the  bottom  of  the  vessel  C. 
A  bell  jar.  or  large  beaker  F  is  supported  over  the  battery  jar,  and  under  its  edge 
extends  the  end  of  a  tube  G  connected  with  a  source  of  hydrogen.  When  hydro- 
gen is  admitted  to  the  space  under  the  cover  F,  it  passes  into  the  porous  jar 
faster  than  the  air  within  passes  out,  developing  a  pressure  within  the  jar.  This 
is  communicated  to  the  surface  of  the  water  in  C,  forcing  some  of  it  out  through 
the  jet  E  as  a  fountain. 

The  meaning  of  laws  in  science.  The  four  laws  just  considered  are 
merely  general  statements  in  regard  to  the  conduct  of  gases  as  deter- 
mined by  experiment.  Like  all  other  scientific  laws,  they  offer  no 


PEOPERTIES  OF  GASES  51 

explanation  of  the  facts  which  they  state,  nor  do  they  place  any  re- 
striction upon  nature  which  compels  obedience,  as  the  laws  enacted 
by  a  legislature  bind  society.  They  are  simply  concise  statements  of 
what  might  be  called  the  habits  of  nature  as  observed  in  experiment. 

Forming  a  theory.  It  is  certainly  a  very  striking  fact  that  all  gas- 
eous substances  behave  in  so  simple  a  manner,  quite  irrespective  of 
their  chemical  nature.  It  would  appear  most  probable  that  this  must 
be  due  to  some  very  simple  mechanical  structure  which  gases  have  in 
common,  and  the  mind  at  once  begins  to  imagine  a  mechanical  model 
which,  if  real,  would  act  in  the  same  manner.  The  process  of  con- 
structing a  mental  picture  of  this  kind  is  called  forming  a  theory. 
The  theory  which  has  proved  to  be  the  most  satisfactory  in  connec- 
tion with  the  properties  of  gases  is  known  as  the  kinetic  theory.  It 
will  be  instructive  to  follow  it  out  to  some  extent,  as  it  will  show 
very  clearly  how  a  theory  is  developed,  and  it  will  be  found  very 
useful  in  subsequent  pages. 

The  chief  points  in  the  kinetic  theory.  A  number  of  points  may  be 
presented  which  must  be  taken  into  account  in  any  theory  which  may 
be  framed  in  regard  to  the  nature  of  gases : 

1.  All  gases  appear  to  have  the  same  mechanical  structure,  since 
they  respond  in  the  same  way  to  energy  changes  such  as  those  of 
temperature  and  pressure. 

2.  They  cannot^  be  in  any  sense  continuous  matter,  but  must  be 
extremely  porous,  since  they  are  always  very  compressible  and  also 
tend  to  expand  indefinitely. 

3.  The  pressure  which  gases  exert  cannot  be  the  thrust  of  a  rigid 
body,  as  of  a  spring,  and,  under  the  circumstances,  almost  the  only 
other  way  in  which  we  can  imagine  the  application  of  a  pressure  is  by 
the  momentum  of  moving  bodies.   Hence  we  may  imagine  a  gas  to  be 
made  up  of  moving  particles  whose  aggregate  impact  is  the  cause  of 
the  pressure  exerted  upon  the  walls  of  a  containing  vessel. 

,^4.  Boyle's  law  states  that  when  the  volume  is  reduced  one  half, 
the  pressure  doubles.  This  is  in  accord  with  the  picture  that  we  are 
drawing,  for  in  the  half  volume  the  particles  will  strike  the  walls  of 
the  vessel  twice  as  often,  and  so  exert  twice  the  force  on  the  same 
area  in  a  given  time. 

5.  Since  rise  of  temperature  increases  the  pressure,  it  must  in  some 
way  increase  the  kinetic  energy  of  the  moving  particles.  This  might 
be  brought  about  by  increasing  either  the  mass  or  the  speed  of  the 


52  *    GENERAL  CHEMISTRY 

particles,  for  their  kinetic  energy  is  equal  to  the  expression  \-  ms2, 
in  which  m  represents  the  mass,  and  *  the  speed.  Experiment  shows 
that  the  mass  of  a  gas  is  not  changed  by  heat ;  so  it  must  be  the  speed 
of  the  particles  which  is  affected. 

6.  The  pressure  of  all  gases  increases  equally  for  equal  rise  in  tem- 
perature ;  so  the  aggregate  energy  of  the  various  kinds  of  gas  particles 
must  increase  to  the  same  degree.    Now  it  can  be  shown  that  the 
masses  of  two  kinds  of  gas  particles  are  different ;  therefore  the  speed 
of  the  particles  must  be  increased  inversely  as  the  square  root  of  their 
mass,  in  order  that  the  value  J-  ms2  may  be  equally  increased  for  two 
different  particles. 

7.  All  the  facts  of  diffusion  are  in  accord  with  the  idea  that  gases 
are  made  up  of  moving  particles,  and  Graham's  law,  which  states  that 
the  rate  of  diffusion  is  inversely  proportional  to  the  square  root  of  the 
density  of  a  gas,  is  but  another  way  of  stating  the  conclusions  of  the 
preceding  paragraph. 

8.  Finally,  our  picture  becomes  clearer  and  more  harmonious  if  we 
assume  that  in  equal  volumes  of  all  gases  there  is  the  same  number 
of  particles.    This  conclusion  may  be  deduced  mathematically  from 
the  quantitative  laws  we  have  had  before  us,  and  it  is  in  accord  with 
many  chemical  facts  to  be  discussed  later  on.    This   assumption  is 
known  as  Avogadro's  hypothesis,  and  the  particles  with  which  it  is 
concerned  have  been  named  molecules. 

Summary  of  the  kinetic  theory.  As  a  picture  which  gives  a  graphic 
representation  of  the  simple  conduct  of  gases,  the  kinetic  theory 
suggests  that  all  gases  are  made  up  of  small  particles  (molecules), 
relatively  far  apart  and  in  motion,  that  equal  volumes  of  all  gases 
contain  the  same  number  of  molecules,  whose  momentum  is  the  cause 
of  pressure,  and  that  the  kinetic  energy  of  all  molecules  is  increased 
equally  by  a  given  rise  in  temperature,  the  increase  being  due  to  the 
increased  speed  of  the  particles. 

Value  of  a  theory.  The  value  of  such  a  theory  is  at  once  apparent. 
It  presents  a  mental  picture  which  assists  the  memory  in  retaining  a 
great  variety  of  facts.  It  suggests  many  experiments  which  otherwise 
might  never  be  undertaken,  for  our  first  impulse,  after  forming  such  a 
theory,  is  to  test  it  experimentally  in  every  possible  way.  It  enables 
us  to  form  a  probable  opinion  in  cases  where  experiment  has  not  yet 
made  a  definite  decision.  It  usually  corrects  errors  which  have  crept 
into  the  body  of  our  knowledge  as  a  result  of  faulty  experiments. 


PROPERTIES  OF  GASES  53 

There  is,  however,  a  real  peril  in  accepting  a  theory.  The  whole 
picture  may  be  wrong,  yet  it  may  seem  so  plausible  that  we  rest 
contented  with  it  and  fail  to  see  its  faults.  Its  very  plausibility  may 
prevent  us  from  making  experiments  which  would  disclose  the  error 
in  the  theory  and  put  us  on  the  right  track. 

All  such  theories  are  best  regarded  as  mere  conveniences.  Doubt- 
less they  express  the  true  nature  of  things  in  many  cases,  but  in  others 
they  do  not.  They  are  useful  as  long  as  we  regard  them  as  con- 
veniences and  as  open  to  constant  revision  and  modification  as  our 
knowledge  grows,  but  they  are  a  real  disadvantage  when  we  come  to 
regard  them  as  the  final  and  unchangeable  truth. 


CHAPTER  V 

WATER 

Historical.  Following  the  discovery  of  hydrogen,  Cavendish  made  a 
careful  study  of  the  properties  of  the  gas.  In  the  course  of  his 
experiments  he  exploded  a  mixture  of  hydrogen  and  air  and  observed 
that  a  small  amount  of  a  dewlike  substance  was  formed.  He  repeated 
the  experiment  a  number  of  times,  on  some  occasions  substituting 
pure  oxygen  for  air,  and  thus  was  able  to  obtain  a  sufficient  amount 
of  the  liquid  to  make  a  study  of  its  properties.  This  liquid  proved  to 
be  pure  water.  Cavendish  did  not  perceive  the  full  meaning  of  his 
discovery,  however,  and  it  remained  for  Lavoisier,  a  few  years  later, 
to  repeat  and  properly  interpret  the  experiments  of  Cavendish.  He 
proved  beyond  doubt  that  the  water  which  Cavendish  had  obtained 
resulted  from  the  union  of  the  hydrogen  and  oxygen,  and  pointed  out 
that  water  must  therefore  be  regarded  as  a  compound  of  these  two 
elements. 

Occurrence.  The  great  abundance  and  wide  distribution  of  water  are 
facts  familiar  to  all.  Vast  areas  of  the  colder  regions  of  the  globe  are 
covered  with  it  in  the  form  of  ice,  while  in  the  liquid  state  it  covers 
about  five  sevenths  of  the  earth's  surface,  reaching  in  some  places  a 
depth  of  nearly  six  miles.  Large  quantities  occur  in  the  soil,  and  as 
a  vapor  it  is  an  essential  constituent  of  the  atmosphere.  It  likewise 
constitutes  more  than  half  the  weight  of  living  organisms.  For 
example,  nearly  70  per  cent  of  the  human  body  is  water.  The  water 
content  of  some  of  the  more  common  foods  is  given  in  the  table 
on  page  308. 

The  composition  of  natural  waters.  All  natural  waters  contain  more 
or  less  foreign  matter,  either  in  solution  or  held  in  suspension.  Even 
the  water  which  falls  to  the  earth  in  the  form  of  rain  contains  parti- 
cles of  dust,  as  well  as  small  quantities  of  gases  absorbed  from  the 
atmosphere.  Upon  reaching  the  earth's  surface  it  dissolves  mineral 
matter  present  in  the  rocks  and  soil,  such  as  common  salt  and  com- 
pounds of  calcium,  magnesium,  and  iron.  Waters  containing  such 
substances  in  solution  are  commonly  spoken  of  as  hard  waters  or,  if 

54 


WATER  55 

large  amounts  of  mineral  matter  are  present,  as  mineral  waters.  The 
quantity  and  nature  of  the  substances  present  vary  with  the  nature 
of  the  rocks  and  soil  with  which  the  water  comes  in  contact.  The 
weight  of  such  matter  present  in  1  1.  of  average  well  water  varies  from 
0.1  to  0.5  g.  Much  larger  quantities  are  present  in  the  waters  from 
some  springs  and  very  deep  wells.  The  waters  of  the  ocean  contain 
over  3.5  per  cent  of  mineral  matter,  more  than  three  fourths  of  which 
is  common  salt.  In  addition  to  mineral  matter  natural  waters  contain 
more  or  less  organic  matter  in  solution  or  held  in  suspension.  This  con- 
sists not  only  of  inanimate  matter,  derived  from  the  decay  of  organic 
bodies  on  the  earth's  surface  or  present  in  sewage,  but  also  of  certain 
forms  of  living  microorganisms  which  usually  accompany  such  prod- 
ucts. Waters  taken  from  shallow  wells  or  streams  in  thickly  populated 
districts  are  likely  to  contain  a  considerable  quantity  of  such  matter. 

Effect  upon  health  of  the  foreign  matter  in  water.  Since  natural 
waters  constitute  the  ordinary  supply  for  drinking  and  household 
purposes,  it  becomes  of  importance  to  inquire  into  the  effect  of  the 
foreign  matter  in  such  waters  when  taken  into  the  system.  Experience 
has  shown  that  the  mineral  matter  commonly  found  in  water  is  not,  as 
a  rule,  injurious  to  health.  In  fact,  the  presence  of  a  certain  amount 
of  such  matter  is  probably  advantageous,  supplying  a  portion  of  the 
mineral  constituents  necessary  for  the  formation  of  the  solid  tissues 
of  the  body. 

As  previously  stated,  the  organic  matter  present  in  water  consists 
of  inanimate  products  as  well  as  of  living  microorganisms.  The 
amount  of  the  former  commonly  present  in  a  water  used  for  drinking 
purposes  is  so  small  that  it  is  practically  without  effect  upon  the 
system.  Of  course,  if  present  in  large  amounts,  or  if  poisonous,  as  in 
the  case  of  sewage,  sickness  would  result  from  its  consumption.  On 
the  other  hand,  the  presence,  in  water,  of  any  considerable  number  of 
microorganisms  renders  it  dangerous  as  a  drinking  water.  It  is  true 
that  many  of  these  organisms  are  without  injurious  effect  upon  the 
system,  but  it  is  likewise  true  that  others  are  the  direct  cause  of 
disease.  Thus  it  is  known  that  a  transmissible  disease  such  as  typhoid 
fever  is  due  to  certain  microorganisms  which  find  entrance  into  the 
body.  It  is  easily  possible  for  these  organisms  to  find  their  way, 
through  sewage,  from  a  person  afflicted  with  the  disease  into  a  poorly 
protected  water  supply,  and  so  contaminate  the  water.  It  is  largely 
in  this  way  that  typhoid  fever  is  spread.  The  general  conclusion  may 


56  GENERAL  CHEMISTRY 

therefore  be  drawn  that,  save  in  exceptional  cases,  any  sickness  trace- 
able to  the  water  supply,  is  due  to  the  presence  in  the  water,  not  of 
mineral  matter  or  even  of  inanimate  organic  matter,  but  to  certain 
living  microorganisms. 

The  detection  of  impurities  in  water.  The  total  amount  of  solid 
matter  present  in  any  given  water  is  easily  determined  by  evaporating 
a  definite  volume  of  the  water  to  dryness  and  weighing  the  residue. 
This  residue  may  then  be  subjected  to  further  investigation  and  the 
nature  of  the  mineral  matter  determined.  A  statement  of  the  mineral 
matter  present  in  a  water,  including  the  percentages  of  each  kind,  is 
commonly  termed  a  mineral  analysis.  Such  an  analysis  is  of  impor- 
tance in  determining  whether  or  not  a  water  is  adapted  for  manufac- 
turing purposes,  such  as  for  use  in  a  steam  boiler.  On  the  other  hand, 
if  one  wishes  to  determine  whether  a  water  is  wholesome  for  drinking, 
a  so-called  sanitary  analysis  is  required.  Such  an  analysis  includes 
primarily  not  only  the  determination  of  the  organic  matter  present 
in  the  water,  but  also  of  the  decomposition  products  formed  by  the 
decay  of  such  matter  (such  as  ammonia,  nitrites,  and  nitrates).  From 
what  has  been  said  it  might  be  inferred  that  a  bacteriological  examina- 
tion alone  would  decide  the  question.  While  it  is  true  that  such  an 
examination  is  of  the  greatest  importance,  yet  it  is  equally  true  that  the 
determination  of  the  inanimate  organic  matter  present,  together  with 
the  products  of.  its  decomposition,  is  of  equal  value  and  supplements 
the  knowledge  gained  from  a  bacteriological  examination ;  for  the 
disease-producing  organisms  find  their  way  into  a  water  supply  through 
the  sewage  or  drains,  and  are  therefore  accompanied  by  other  organic 
matter,  the  presence  of  which  in  a  water  supply  at  once  indicates  pol- 
lution. Such  a  water  should  therefore  not  be  used,  for,  although  it 
may  be  temporarily  free  from  disease-producing  organisms,  the  condi- 
tions are  such  that  their  introduction  may  take  place  at  any  time. 

It  may  be  added  that  the  physical  properties  of  a  drinking  water 
rarely  give  any  conclusive  evidence  as  to  its  purity.  A  water  may  be 
unfit  for  drinking  and  yet  be  perfectly  clear  and  odorless.  Neither 
can  any  reliance  be  placed  on  the  simple  methods  sometimes  given  for 
testing  the  purity  of  water.  Only  the  trained  chemist  and  bacteri- 
ologist can  carry  out  such  methods  of  analysis  as  are  trustworthy. 

The  removal  of  foreign  matter  from  natural  waters  ;  distilled  water. 
Inasmuch  as  all  natural  waters  contain  foreign  matter,  it  becomes  of 
interest  to  inquire  into  the  methods  used  for  removing  such  matter 


WATER 


57 


and  so  for  obtaining  chemically  pure  water.  The  process  employed 
for  this  purpose  is  known  as  distillation,  and  consists  in  boiling  the 
water  and  condensing  the  resulting  steam. 

As  commonly  carried  out  on  a  small  scale  in  the  laboratory  the  process 
is  as  follows :  The  sample  of  water  is  poured  into  the  flask  A  (Fig.  27)  and 
boiled.  The  resulting  steam  is  conducted  through  the  condenser  B,  which 
usually  consists  "of 
a  narrow  glass  tube 
sealed  within  a  larger 
one.  A  current  of 
cold  water  which 
is  admitted  at  C 
and  escapes  at  D  is 
continuously  passed 
through  the  space  be- 
tween the  two  tubes. 
The  inner  tube  is 
thus  kept  cool,  and 
the  steam,  in  passing 
through  it,  is  con- 
densed and  collects 
in  E.  FIG.  27 

The  water  formed  by  the  condensation  of  steam  is  known  as  dis- 
tilled water.  The  mineral  matter  present  in  the  original  water,  being 
nonvolatile,  remains  in  the  container  in  which  the  water  is  being 
boiled.  The  organic  matter  is  also  largely  left  in  the  container.  A 
small  amount  of  it,  however,  may  be  decomposed  into  volatile  products, 
in  which  case  these  would  pass  over  with  the  steam  and  be  present  in 
the  distilled  water.  The  percentage  of  such  matter  present  in  distilled 
water  is  so  small,  however,  that  it  is  practically  without  effect  in 
the  chemical  processes  in  which  pure  water  is  employed,  except  in 
a  very  few  cases  where  extreme  purity  is  required.  If  it  is  desired  to 
remove  such  traces  of  foreign  matter,  the  distilled  water  is  treated 
with  certain  reagents  which  combine  with  the  impurities  to  form 
nonvolatile  compounds,  and  the  process  of  distillation  is  repeated. 
In  such  cases  the  steam  is  condensed  in  a  tube  made  of  tin  rather 
than  of  glass,  since  this  metal  is  more  resistant  than  glass  to  the 
action  of  steam.  For  the  distillation  of  water  on  a  large  scale  the 
water  is  heated  in  a  boiler  made  of  copper  or  iron,  and  the  steam  is 
condensed  in  a  tube  made  of  tin  wound  into  the  form  of  a  spiral  and 
surrounded  by  cold  water. 


58  GENERAL  CHEMISTRY 

The  purification  of  water  for  sanitary  purposes.  By  the  term  pure 
water  as  commonly  used  we  do  not  necessarily  mean  water  from  which 
all  foreign  matter  has  been  removed,  but  rather  one  in  which  any  such 
matter  present  is  not  injurious  to  health.  Such  objectionable  matter 
may  of  course  be  removed  by  the  process  of  distillation,  as  described 
above,  but  water  may  also  be  purified  by  the  process  of  boiling  or  by 
filtration. 

Effect  of  boiling.  The  statement  has  been  made  that  it  is  the  living 
microorganisms  present  in  water  which  render  it  unwholesome.  These 
organisms  cannot  withstand  a  high  temperature.  If  a  water  is  boiled 
a  few  minutes,  therefore,  any  organisms  present  are  killed  and  the 
water  is  rendered  safe  for  drinking.  The  effect  of  boiling  is  not  to 
remove  the  microorganisms  but  simply  to  destroy  their  vitality.  It 
will  be  shown  in  a  subsequent  chapter  that  some  compounds  of  cal- 
cium and  magnesium,  if  present  in  the  water,  are  thrown  out  of 
solution  by  the  process  of  boiling;  otherwise  the  mineral  content  is 
not  effected. 

Effect  of  filtration.  In  the  process  of  filtration  the  water  is  passed 
through  some  medium,  such  as  charcoal,  which  possesses  the  property 
of  absorbing  any  organic  matter  present  in  the  water.  To  be  effective 
such  a  filter  must  be  kept  clean,  since  it  is  evident  that  charcoal  is  use- 
less as  a  filtration  medium  after  its  pores  become  filled  with  impuri- 
ties. In  fact,  unless  the  charcoal  is  renewed  from  time  to  time  it  may 
become  so  contaminated  with  microorganisms  that  it  will  serve  as  a 
source  of  pollution  of  the  water  rather  than  of  purification.  A  more 
effective  type  of  filter  is  the  so-called  Chamberlain-Pasteur  filter. 
This  consists  of  one  or  more  cylindrical  cups,  the  pores  of  which 
are  very  minute.  When  attached  to  the  water  faucet  and  the  water 
turned  on,  the  liquid  is  forced  through  these  minute  pores,  while 
any  microorganisms  present  are  strained  out  and  remain  in  the  in- 
terior of  the  cylinder.  Such  a  filter  likewise  removes  any  mineral 
matter  held  in  suspension,  but  has  no  effect  upon  such  matter  that 
is  in  solution. 

The  filtration  of  water  on  a  large  scale.  Many  cities  find  it  neces- 
sary to  take  their  water  supply  from  rivers  and  reservoirs  in  which  the 
water  is  more  or  less  contaminated  with  organic  matter.  Such  a  water 
supply  is  a  source  of  constant  menace  to  the  health  of  the  city.  It 
becomes  necessary,  therefore,  to  find  some  way  of  filtering  the  water 
effectively  on  a  large  scale.  This  can  be  done  by  allowing  the  water 


WATER 


59 


to  pass  through  large  filtration  beds  prepared  from  sand  and  gravel. 
Some  of  the  impurities  are  strained  out  by  the  filter,  while  others  are 
decomposed  by  the  action  of  certain  kinds  of  microorganisms  which 
collect  in  a  gelatinous  layer  on  the  surface  of  the  filter. 

Fig.  28  shows  a  cross  section  of  such  a  filter  bed.  The  water  filters  through 
the  sand  and  gravel  and  passes  into  the  porous  pipe  A,  from  which  it  is  pumped 
into  the  city  mains.  The  fil- 
ters are  usually  covered,  to 
prevent  the  water  from  freez- 
ing in  cold  weather. 


FIG. 28 


In  some  cases  the  water 
before  filtration  is  pumped 
into  large  tanks  and  treated 
with  certain  compounds, 
such  as  alum,  which  form 
in  the  water  a  small  amount 
of  gelatinous  solid.  This 
slowly  settles  to  the  bot- 
tom of  the  tank,  carrying 
with  it  much  of  the  organic  matter  present.  The  effect  of  the  filtration 
of  the  water  supply  upon  the  health  of  a  city  is  shown  by  the  fact 
that  in  general  the  number  of  cases  of  typhoid  fever  in  cities  which 
have  introduced  an  effective  water  purification  system  has  been 
decreased  about  75  per  cent. 

Self-purification  of  water.  It  has  long  been  known  that  water  con- 
taminated with  organic  matter  tends  to  purify  itself  when  exposed  to 
the  air.  For  example,  the  water  in  a  river  made  very  foul  by  the  in- 
troduction of  sewage  at  some  point  gradually  becomes  purer  as  it 
flows  down  the  bed  of  the  river.  While  purification  may  be  influenced 
by  several  causes,  it  is  primarily  due  to  the  oxidation  of  the  organic 
matter  by  oxygen  absorbed  from  the  air.  Streams  of  purer  water  flow- 
ing in  greatly  assist  not  only  by  dilution  but  also  by  bringing  in  addi- 
tional oxygen.  As  the  inanimate  organic  matter  is  thus  gradually 
removed,  many  of  the  microorganisms  die  from  lack  of  nutriment  or 
from  change  in  conditions,  while  others  gradually  settle  to  the  bottom 
of  the  river,  especially  if  the  current  is  slow.  This  method  of  puri- 
fication, however,  cannot  be  relied  upon,  especially  in  thickly  populated 
districts,  to  purify  contaminated  water  so  as  to  render  it  safe  for 
drinking  purposes. 


60 


GENERAL  CHEMISTRY 


While  all  microorganisms  are  destroyed  by  heat,  low  temperature 
has  but  little  effect  upon  many  of  them.  It  follows  that  ice  frozen 
from  impure  water  may  be  contaminated  and  should  not  be  used 
for  household  purposes. 

Properties.  At  ordinary  temperatures  pure  water  is  a  clear,  trans- 
parent liquid.  It  has  a  slightly  bluish  color,  but  this  is  only  noticeable 
in  water  of  considerable  depth,  especially  in  lakes  in  which  the  water 
is  fairly  pure  and  deep.  It  solidifies  at  0°  and  boils  at  100°  under 
the  normal  pressure.  The  density  of  water  varies  with  its  temperature, 
reaching  its  maximum  at  4°,  as  shown  in  the  following  table : 


TEMPERATURE 

DENSITY 

TEMPERATURE 

DENSITY 

0° 

0.9998 

40° 

0.9923 

40 

1.0000 

60° 

0.9833 

10° 

0.9997 

80° 

0.9719 

20 

0.9982 

100° 

0.9586 

Water  possesses  to  a  remarkable  extent  the  property  of  dissolving 
other  substances ;  indeed,  a  greater  variety  of  matter  is  dissolved  by 
water  than  by  any  other  known  solvent.  Many  substances,  such  as 
glass  and  various  kinds  of  rock,  which  are  ordinarily  considered  in- 
soluble in  water,  do,  however,  dissolve  to  a  very  limited  extent. 

The  properties  of  water,  as  well  as  the  ease  with  which  the  pure 
liquid  can  be  obtained,  render  it  well  adapted  for  use  in  the  choice 
of  certain  units.  Thus,  its  boiling  and  freezing  temperatures  are  the 
fixed  points  used  in  the  graduation  of  the  thermometric  scales.  Its 
connection  with  the  unit  of  mass,  as  well  as  of  heat,  has  been  referred 
to  in  Chapter  I.  The  change  of  water  from  any  one  of  the  three 
states,  solid,  liquid,  and  gaseous,  into  any  other  of  these  states  is  al- 
ways attended  by  an  absorption  or  liberation  of  energy  in  the  form 
of  heat.  Such  energy  transformations  always  accompany  changes  in 
state,  not  only  of  water  but  of  matter  in  general.  They  will  there- 
fore be  discussed  under  a  single  head  in  the  following  chapter. 

Chemical  conduct.  A  knowledge  of  the  chemical  conduct  of  water 
is  of  fundamental  importance  for  an  understanding  of  many  chemical 
processes.  The  main  topics  to  be  considered  under  this  head  are  the 
effect  of  heat  upon  water,  the  reaction  between  water  and  certain  ele- 
ments and  compounds,  and  the  part  that  water  plays  in  promoting 
chemical  changes. 


WATER 


61 


1.  The  effect  of  heat  upon  water.  We  have  already  seen  that  it  is 
possible  to  decompose  water  either  by  the  action  of  the  electric  cur- 
rent or  by  passing  its  vapor  over  some  highly  heated  metal  such  as 
iron.  It  is  also  of  interest  to  inquire  whether  the  decomposition  of 
water  into  hydrogen  and  oxygen  can  be  brought  about  by  heat  alone. 
Experiments  have  shown  that  such  a  decomposition  does  occur,  but 
only  to  a  slight  extent,  even  at  very  high  temperatures.  The  fol- 
lowing table  by  Langmuir  gives  the  percentages  of  the  total  quan- 
tity of  water  decomposed  when  heated  to  the  temperature  indicated. 
The  results  must  be  considered  as  only  approximate  ones,,  since  the 
experiments  are  very  difficult  to  carry  out: 


TEMPERATURE 

PERCENTAGE  OF 
WATER  DECOMPOSED 

TEMPERATURE 

PERCENTAGE  OF 
WATER  DECOMPOSED 

1327° 

0.0446 

1727° 

0.504 

1427° 

0.0920 

1927° 

1.21 

1527° 

0.17 

2227° 

3.38 

1627° 

0.302 

2727° 

11.10 

Compounds,  like  water,  which  are  not  readily  decomposed,  especially 
by  heat,  are  commonly  spoken  of  as  stable  compounds. 

When  water  vapor  is  heated  in  a  closed  vessel  to  a  temperature  above  the 
point  at  which  decomposition  begins,  the  percentage  of  the  total  vapor  which  is 
decomposed  gradually  increases  with  the  rise  in  temperature.  If,  however,  the 
temperature  is  maintained  constant  at  any  point,  the  decomposition  apparently 
ceases.  In  reality,  however,  it  is  known  that  the  decomposition  continues,  but 
that  the  quantity  of  vapor  decomposed  is  exactly  equal  to  that  formed  again  by 
the  union  of  the  oxygen  and  hydrogen.  In  other  words,  the  water  vapor  is  in 
equilibrium  with  its  decomposition  products,  namely,  oxygen  and  hydrogen.  The 
change,  therefore,  like  the  change  of  oxygen  into  ozone  and  the  reaction  between 
iron  and  steam,  is  a  reversible  one  and  may  be  expressed  as  follows : 

water  vapor  ^     ^  hydrogen  +  oxygen 

2.  The  action  of  water  upon  certain  elements.    Reference  has  been 
made  to  the  fact  that  some  of  the  elements,  such  as  iron,  react  with 
water  under  proper  conditions,  combining  with  the  oxygen  and  thus 
liberating  the  hydrogen.    On  the  other  hand,  a  few  of  the  elements, 
such  as  fluorine  and  chlorine,  have  just  the  opposite  action,  combining 
with  the  hydrogen  and  liberating  the  oxygen. 

3.  The  action  of  water  upon  oxides.    Water  combines  with  many  of 
the  oxides  to  form  important  compounds.    It  is  convenient  to  divide 
these  oxides  into  two  general  classes,  according  to  the  nature  of  the 


62  GENERAL  CHEMISTRY 

compounds  resulting  from  their  union  with  water :  (1)  The  members 
of  the  one  class  combine  with  water  to  form  acids.  This  class 
is  represented  by  such  oxides  as  sulfur  dioxide  and  phosphorus 
pentoxide,  compounds  formed  by  the  combustion  of  sulfur  and  phos- 
phorus respectively.  (2)  The  members  of  the  second  class  of  oxides, 
on  the  other  hand,  combine  with  water  to  form  compounds  known 
as  bases,  to  which  reference  has  already  been  made  (p.  38).  To 
this  class  belong,  for  example,  the  oxide  of  calcium  (ordinary  lime) 
which  combines  with  water  to  form  the  base  known  as  calcium 
hydroxide,  or  slaked  lime.  These  two  classes  of  compounds,  namely, 
acids  and  bases,  are  of  the  greatest  importance,  and  their  properties 
will  be  discussed  in  detail  in  a  later  chapter.  It  is  sufficient  at  this 
time  to  note  that  in  many  respects  they  are  opposite  in  character, 
and  that  they  react  with  each  other  to  form  water  and  compounds 
called  salts. 

4.  The  action  of  water  upon  compounds  other  than  oxides.  When  crys- 
tals of  ordinary  alum  are  heated,  water  is  evolved,  while  the  crystals 
crumble  to  a  fine  powder  commonly  known  as  burnt  alum.  When 
this  burnt  alum  is  dissolved  in  water  and  the  resulting  solution  allowed 
to  evaporate,  there  is  deposited  a  crystalline  compound,  formed  by  the 
union  of  the  burnt  alum  with  water,  which  is  identical  in  composition 
with  the  original  crystals.  Similarly,  when  the  blue  crystalline  com- 
pound known  as  copper  sulfate  or  blue  vitriol  is  heated,  water  is 
evolved,  the  blue  color  fading  at  the  same  time  until  there  remains 
a  white,  powdery  residue.  As  in  the  case  of  alum,  the  blue  crystals 
may  again  be  formed  from  this  residue  by  dissolving  it  in  water  and 
evaporating  the  solution.  Many  other  compounds,  especially  those 
belonging  to  the  class  known  as  salts,  act  like  alum  and  copper  sul- 
fate, evolving  water  when  heated,  but  again  combining  with  it  in  solu- 
tion to  form  the  original  compound.  The  compounds  which  combine 
with  the  water  are  commonly  spoken  of  as  anhydrous  substances,  while 
the  compounds  formed  by  their  union  with  water  are  termed  hydrates. 
Thus  the  blue  crystalline  compound  referred  to  above  is  a  hydrate  of 
copper  sulfate,  while  the  white  residue  formed  on  heating  this  hydrate 
is  the  anhydrous  copper  sulfate.  These  hydrates  are  true  chemical 
compounds,  any  given  hydrate  being  formed  by  the  union  of  definite 
weights  of  the  anhydrous  substance  and  water.  Thus  any  given 
weight  of  anhydrous  copper  sulfate  always  combines  with  56.43  per 
cent  of  its  weight  of  water  to  form  the  ordinary  hydrate.  Many 


WATER  63 

anhydrous  substances,  however,  combine  with  different  percentages  of 
water  to  form  different  hydrates. 

The  hydrates  are  not,  as  a  rule,  very  stable,  but  may  be  decomposed 
into  the  constituents  from  which  they  are  formed,  namely,  water  and 
the  anhydrous  substances.  As  in  the  examples  given  above,  this  de- 
composition can  be  brought  about  by  heat,  the  temperature  at  which 
the  decomposition  takes  place  varying  with  the  nature  of  the  hydrate. 
In  some  cases  it  will  even  take  place  at  ordinary  temperatures.  Thus 
the  clear  crystals  of  the  hydrate  of  sodium  sulfate,  when  exposed  to 
dry  air,  gradually  lose  water,  leaving  an  opaque  mass  of  the  anhy- 
drous sulfate.  Such  substances  are  said  to  be  efflorescent.  The  gen- 
eral principles  underlying  the  decomposition  of  the  hydrates  will  be 
discussed  under  the  subject  of  equilibrium. 

The  water  which  combines  with  compounds  to  form  hydrates  is 
commonly  termed  water  of  crystallization,  for  the  reason  that  the 
hydrates  are  crystalline  in  character  and,  when  heated,  lose  water, 
with  an  accompanying  loss  of  crystalline  structure. 

It  must  not  be  supposed  that  all  crystalline  substances  contain  water  of  crys- 
tallization. The  majority  of  minerals,  such  as  quartz  and  the  diamond,  are 
crystalline  without  having  water  of  crystallization.  Likewise  some  salts,  such  as 
sodium  chloride,  crystallize  from  water  in  the  anhydrous  form. 

Substances  crystallizing  from  water  may  contain  varying  amounts  of  water 
simply  inclosed  mechanically  in  the  crystal.  When  such  substances  are  heated, 
this  water  is  changed  into  steam,  and  the  crystal  is  thereby  torn  apart,  often 
with  a  crackling  sound.  In  such  cases  the  substance  is  said  to  decrepitate.  This 
property  is  especially  noticeable  in  crystals  of  common  salt. 

The  term  water  of  crystallization  seems  to  imply  the  presence  of  water  as 
such  in  the  crystals.  There  is  no  evidence,  however,  that  this  is  so.  That  a  com- 
pound evolves  water  on  being  heated  is  not  a  proof  that  the  water  as  such  is 
present  in  it,  since  the  compound  may  contain  hydrogen  and  oxygen  which 
unite  to  form  water  in  the  process  of  decomposition. 

Water  may  also  react  with  many  compounds  in  other  ways  than  by 
direct  combination.  It  was  stated  above  that  acids  and  bases  react 
with  each  other  to  form  water  and  compounds  known  as  salts.  In  the 
case  of  some  salts  this  reaction  is  reversible  and  may  therefore  be 
expressed  as  follows : 

acid  4-  base  <    >  salt  -f-  water 

If,  therefore,  water  is  added  to  such  a  salt,  at  least  a  portion  of  the  salt 
reacts  with  the  water  to  form  an  acid  and  a  base.  This  reaction  is 
commonly  termed  hydrolysis,  and  will  be  fully  discussed  later. 


64  GENERAL  CHEMISTRY 

5.  The  action  of  water  in  promoting  chemical  changes.  Many  substances 
which  have  no  action  upon  each  other  in  the  absence  of  water  readily 
enter  into  combination  in  its  presence.  The  reason  for  this  is  not 
always  clear.  In  some  cases  the  increased  activity  seems  to  be  due 
entirely  to  the  fact  that,  when  in  solution,  the  substances  are  brought 
into  more  intimate  contact.  On  the  other  hand,  all  acids,  bases,  and 
salts,  when  dissolved  in  water,  undergo  at  least  a  partial  decompo- 
sition, and  the  greater  activity  of  these  compounds  when  in  aqueous 
solution  is  connected  with  this  fact  (see  p.  148).  Attention  has  already 
been  called  to  the  marked  catalytic  effect  of  moisture.  Even  a  trace 
of  it  often  has  a  marked  influence  in  increasing  the  speed  of  certain 
chemical  changes. 

The  determination  of  the  exact  composition  of  water.  Many  very 
careful  experiments  have  been  made  for  the  purpose  of  determining, 
with  as  great  accuracy  as  possible,  the  ratio  in  which  hydrogen  and 
oxygen  are  present  in  water,  and  it  is  worth  our  while  to  study  some- 
what in  detail  the  methods  which  have  been  employed,  since  they 
serve  to  illustrate  in  a  general  way  the  methods  used  in  determining 
the  composition  of  other  compounds. 

Two  general  methods  of  procedure  are  available  for  determining 
the  composition  of  a  compound :  first,  the  method  of  analysis,  in  which 
a  given  weight  of  the  compound  is  separated  either  directly  or  indi- 
rectly into  its  constituent  elements  and  the  identity  and  weight  of  each 
determined;  second,  the  method  of  synthesis,  which  consists  in  deter- 
mining the  proportion  in  which  the  constituent  elements  unite  to  form 
the  compound,  and  which  is  therefore  just  the  opposite  of  analysis. 

1.  Methods  based  on  analysis.  It  will  be  recalled  that  water  may  be 
easily  decomposed  into  its  constituents  by  the  electric  current.  It 
would  naturally  seem  that  the  exact  composition  of  water  could  easily 
be  determined  in  this  way,  since  the  volumes  of  the  gases  liberated  can 
readily  be  measured  with  accuracy,  and  if  we  know  their  densities,  ,the 
weights  of  the  gases  so  liberated  can  be  calculated.  When  the  ex- 
periment is  carried  out,  however,  the  results  obtained  are  not  con- 
cordant, although  in  general  the  volume  of  the  hydrogen  liberated  is 
slightly  more  than  double  the  volume  of  the  oxygen.  Experiments 
prove  that  the  method  is  subject  to  several  sources  of  error.  For  ex- 
ample, a  portion  of  the  oxygen  liberated  is  converted  into  ozone. 
Moreover,  the  water  through  which  the  liberated  gases  bubble  (see 
Fig.  5)  dissolves  more  of  the  oxygen  than  of  the  hydrogen.  The 


WATER 


65 


ratio  between  the  amounts  of  hydrogen  and  oxygen  obtained  in  this 
process,  therefore,  does  not  represent  with  great  accuracy  the  ratio  in 
which  they  are  combined  in  water.  More  accurate  results  are  obtained 
by  the  synthetic  methods  described  in  the  succeeding  pages. 

2.  Methods  based  on  synthesis.  In  the  synthetic  methods  we  deter- 
mine the  quantities  of  oxygen  and  hydrogen  which  combine  directly 
to  form  water ;  or  we  may  determine  the  quantity  of  either  of  the 
elements  which  enter  into  the  combination,  and  then  determine  the 
weight  of  the  resulting  water,  the  difference  * 
between  these  two  weights  being  equal  to  the 
weight  of  the  other  element  entering  into  com- 
bination. Three  modifications  of  the  method 
will  be  described : 

(a)  Method  used  in  lecture  room,.  A  descrip- 
tion of  the  method  as  commonly  carried  out  for 
purposes  of  illustration  in  the  lecture  room  will 
serve  to  show  the  general  principle  involved. 

In  this  method  the  volumes  of  hydrogen  and 
oxygen  combining  to  form  water  are  directly  deter- 
mined. The  combination  of  the  two  gases  is  brought 
about  in  a  tube  called  a  eudiometer.  This  is  a  grad- 
uated tube  about  60  cm.  long  and  2  cm.  wide,  closed 
at  one  end  (Fig.  29).  Xear  the  closed  end  two  plati- 
num wires  are  fused  through  the  glass,  the  ends  of  the  wires  within  the  tube 
being  separated  by  a  space  of  2  or  3  mm.  The  tube  is  entirely  filled  with  mer- 
cury and  inverted  in  a  vessel  of  the  same  liquid.  Pure  hydrogen  is  passed  into 
the  tube  until  it  is  about  one  fourth  filled.  The  volume  of  the  gas  is  then  read 
off  on  the  scale  and  reduced  to  standard  conditions.  An  approximately  equal 
volume  of  pure  oxygen  is  then  introduced  and  the  volume  again  read  off  and 
reduced  to  standard  conditions.  This  gives  the  total  volume  of  the  two  gases. 
From  this  the  volume  of  the  oxygen  introduced  may  be  determined  by  subtract- 
ing from  it  the  volume  of  the  hydrogen.  The  combination  of  the  two  gases  is 
now  brought  about  by  connecting  the  two  platinum  wires  with  an  induction  coil 
and  passing  a  spark  from  one  wire  to  the  other.  Immediately  a  slight  explosion 
occurs.  The  mercury  in  the  tube  is  at  first  depressed  because  of  the  expansion 
of  the  gases  due  to  the  heat  generated  in  the  reaction,  but  at  once  rebounds, 
taking  the  place  of  the  gases  which  have  combined  to  form  water.  The  volume 
of  the  water  in  the  liquid  state  is  so  small  that  it  may  be  disregarded  in  the 
calculations.  In  order  that  the  temperature  of  the  residual  gas  and  the  mercury 
may  become  uniform,  the  apparatus  is  allowed  to  stand  for  a  few  minutes.  The 
volume  of  the  gas  is  then  read  off  and  reduced  to  standard  conditions,  so  that 
it  may  be  compared  with  the  volumes  of  the  hydrogen  and  oxygen  originally 
taken.  The  residual  gas  is  then  tested  in  order  to  ascertain  whether  it  is  hydro- 
gen or  oxygen,  experiments  having  proved  that  it  is  never  a  mixture  of  the  two. 


FIG.  29 


66 


GENERAL  CHEMISTEY 


From  the  information  thus  obtained  the  composition  of  water  may  be  calculated. 
Thus,  suppose  the  readings  wef e  as  follows  : 

Volume  of  hydrogen  taken 20.3  cc. 

Volume  of  hydrogen  and  oxygen 38.7  cc. 

Volume  of  oxygen 18.4  cc. 

Volume  of  gas  left  after  combination  has  taken  place  (oxygen)    8.3  cc. 

The  20.3  cc.  of  hydrogen  have  combined  with  18.4  cc.  minus  8.3  cc.  (or  10.1  cc.) 
of  oxygen,  or  approximately  2  volumes  of  hydrogen  have  combined  with  1  of 
oxygen.    Since  oxygen  is  15.88  times  as  heavy  as  hydrogen,  the  proportion  by 
weight  in  which  the  two  ga*ses  combined  is  1  part  of  hydrogen  to  7.94  of  oxygen. 
Precaution.  If  the  two  gases  are  introduced  into  the  eudiometer  in  the  exact 
proportions  in  which  they  combine,  after  the  combination  has  taken  place  the 
liquid  will  rise  and  completely  fill  the  tube.    Under  these  conditions,  however, 
the  tube  will  be  broken  by  the  sudden  upward  rush 
of  the  liquid.    Hence,  in  performing  the  experiment 
care  is  taken  to  introduce  an  excess  of  one  of  the 
gases  to  serve  as  a  cushion. 

A  more  convenient  form  of  eudiometer.  A  form  of 
eudiometer  (Fig.  30)  different  from  that  shown  on 
page  65  is  sometimes  used  to  avoid  the  calculations 
necessary  in  reducing  the  volumes  of  the  gases  to 
the  same  conditions  of  temperature  and  pressure  in 
order  to  make  comparisons.  With  this  apparatus  it 
is  possible  to  take  the  readings  of  the  volumes  under 
the  same  conditions  of  temperature  and  pressure,  and 
thus  compare  them  directly.  The  apparatus  is  filled 
with  mercury  and  the  gases  introduced  into  the  tube 
A.  The  experiment  is  carried  out  as  in  the  preceding 
one,  except  that,  before  taking  the  reading  of  the  gas 
volumes,  mercury  is  either  added  to  the  tube  B  or 
withdrawn  from  it  by  means  of  the  stopcock  C,  until 
it  stands  at  exactly  the  same  height  in  both  tubes. 
The  gas  inclosed  in  tube  A  is  then  under  atmospheric 
pressure.  The  temperature  of  the  gas,  as  well  as  the 
pressure  io  which  it  is  subjected,  being  the  same  at  the 

conclusion  of  the  experiment  as  at  the  beginning,  the  volumes  of  the  hydrogen  and 
oxygen  and  of  the  residual  gas  may  be  directly  compared  as  read  off  from  the  tube. 

(b)  Method  used  by  Berzelius  and  Dumas.  The  work  of  these  inves- 
tigators is  of  interest  from  a  historical  standpoint,  since  they  were  the 
first  to  determine  the  composition  of  water  with  any  great  accuracy. 
The  method  used  is  a  very  ingenious  one,  the  weights  of  the  hydro- 
gen and  oxygen  being  determined  by  indirect  methods  and  not  by 
direct  weighing  of  the  gases,  which  is  not  easily  done.  The  method 
was  first  used  by  Berzelius  in  1820  and  later  in  1843,  with  greater 
refinement,  by  Dumas. 


FIG.  30 


WATER 


6T 


Fig.  31  will  serve  to  illustrate  the  method  used  by  Dumas.  Hydrogen  is  gen- 
erated in  A  by  the  action  of  dilute  sulfuric  acid  on  zinc.  It  is  conducted  from 
the  generator  through  the  tubes  B  and  C,  which  are  filled  with  different  sub- 
stances designed  to  remove  all  possible  impurities  from  the  gas.  The  pure  gas 
is  then  conducted  into  the  bulb  D,  which  is  partly  filled  with  pure  copper  oxide 
and  kept  at  a  high  temperature.  A  portion  of  the  hydrogen  combines  with 
oxygen  taken  from  the  copper  oxide  to  form  water,  most  of  which  is  condensed 
in  E.  The  remainder  is  absorbed  by  the  substance  contained  in  the  tubes  F 
and  G.  H  represents  a  tube  filled  with  some  substance  which  prevents  any 


FIG.  31 

moisture  entering  the  tube  G  from  the  air.  The  weight  of  the  water  formed  is 
determined  by  noting  the  increase  in  weight  of  the  bulb  E,  as  well  as  of  the 
tubes  F  and  G.  The  weight  of  the  oxygen  entering  into  combination  is  deter- 
mined by  noting  the  loss  in  weight  of  the  bulb  D,  since  this  loss  is  due  solely 
to  the  oxygen  taken  up  from  the  copper  oxide  by  the  hydrogen  to  form  water. 
The  difference  between  the  weight  of  the  water  formed  and  that  of  the  oxygen 
in  the  water  represents  the  weight  of  the  hydrogen.  Dumas  carried  out  this 
experiment  nineteen  times.  He  found  that  the  total  weight  of  water  formed  in 
his  experiments,  together  with  the  weights  of  oxygen  and  hydrogen  present  in 
the  water,  was  as  follows  : 

Total  water  formed 945.439  g. 

Total  oxygen 840.161  g. 

Total  hydrogen 105.278  g. 

The  relative  weights  of  hydrogen  and  oxygen  in  water,  as  determined  by  Dumas, 
are  therefore  105.278  :  840.161,  or  1  part  by  weight  of  hydrogen  to  7.98  of  oxygen. 

(c)  Method  used  by  Morley.  The  method  used  by  Morley  consists  in 
preparing  pure  hydrogen  and  oxygen, -and  determining  not  only  the 
weights  of  the  two  gases  which  combine  to  form  water  but  also  the 
weight  of  the  resulting  water. 

Fig.  32  represents  the  form  of  apparatus  used  by  Morley  for  effecting  the 
combination  of  hydrogen  and  oxygen  and  for  weighing  the  resulting  water.  Ex- 
traordinary precautions  were  taken  to  insure  pure  materials  and  to  eliminate  all 
known  sources  of  error.  The  air  was  first  removed  from  the  apparatus,  which  was 
then  sealed  and  weighed.  A  tube  containing  hydrogen  absorbed  in  palladium 


68 


GENERAL  CHEMISTRY 


was  weighed  and  joined  to  the  apparatus  at  A,  while  a  large  globe  filled  with 
oxygen  was  weighed  and  connected  at  B.  The  two  gases  were  then  admitted  to 
the  apparatus  through  the  tubes  C,  C,  and  their  union  effected  as  they  entered 
by  electric  sparks  passed  between  the  points  at  D.  In  order  that  the  resulting 
steam  might  be  condensed  as  fast  as  formed,  the  apparatus  was  immersed  in  cold 
water  during  the  experiment.  After  from  30  to  35  g.  of  water  had  thus  been 
formed,  the  vessels  from  which  the  hydrogen  and  oxygen 
were  supplied  were  disconnected  and  again  weighed  to  deter- 
mine the  exact  weights  of  hydrogen  and  oxygen  admitted 
to  the  apparatus.  There  still  remained  in  the  apparatus, 
however,  a  certain  amount  of  uncombined  hydrogen  and 
oxygen  which  had  to  be  determined.  To  do  this  the  entire 
apparatus  was  immersed  in  a  freezing  mixture  until  the 
water  which  had  been  formed  by  the  union  of  the  hydrogen 
and  oxygen  was  frozen.  The  mixture  of  the  uncombined 
gases  was  then  withdrawn  through  A  and  B.  Any  moisture 
present  in  the  mixture  was  removed  as  the  gas  passed 
through  the  tubes  E,  E,  which  were  filled  with  phosphorus 
pentoxide,  a  substance  which  has  a  strong  affinity  for  water. 
The  weights  of  hydrogen  and  oxygen  present  in  the  mix- 
ture were  then  determined  by  analysis.  Finally  the  appa- 
ratus itself  was  weighed,  the  increase  in  weight  representing 
the  water  formed  in  the  experiment. 

As  the  average  of  twelve  experiments,  Morley 
found  that  the  proportion  in  which  hydrogen  and 
oxygen  unite  to  form  water  is  as  follows :  by 
weight,  1  part  of  hydrogen  to  7.94  parts  of  oxygen  ; 
by  volume,  2.0024  parts  of  hydrogen  to  1  part  of 
oxygen. 

Comparison  of  results  obtained.  From  the  above 
discussion  it  is  easy  to  see  that  it  is  only  by  experi- 
ment that  the  composition  of  a  compound  can  be  determined.  Differ- 
ent methods  may  lead  to  slightly  different  results.  The  more  accurate 
the  method  chosen,  and  the  greater  the  skill  with  which  the  experiment 
is  carried  out,  the  more  accurate  will  be  the  results.  It  is  universally 
conceded  that  Morley's  results  are  the  most  trustworthy  yet  obtained. 
Relation  between  any  given  volume  of  aqueous  vapor  and  the  vol- 
umes of  the  hydrogen  and  oxygen  which  combine  to  form  it.  When  the 
quantitative  synthesis  of  water  is  carried  out  at  ordinary  temperatures, 
the  water  vapor  formed  by  the  union  of  the  hydrogen  and  oxygen 
at  once  condenses.  The  volume  of  the  resulting  liquid  is  so  small 
that  it  may  be  disregarded  in  making  the  calculations.  If,  however, 
the  experiment  is  carried  out  at  a  temperature  of  100°  or  above,  the 


FIG.  32 


WATER 


69 


water  vapor  formed  is  not  condensed,  and  it  thus  becomes  possible 
to  compare  the  volume  of  the  vapor  with  the  volumes  of  hydrogen 
and  oxygen  which  combined  to  form  it.  In  this  way  it  has  been 
proved  that  2  volumes  of  hydrogen  and  1  volume  of  oxygen  combine 
to  form  exactly  2  volumes  of  water  vapor,  the  volumes  all  being 
measured  under  the  same  conditions  of  temperature  and  pressure.  It 
will  be  noted  that  the  relation  between  these  volumes  may  be  ex- 
pressed by  whole  numbers.  The  significance  of  this  very  important 
fact  will  be  discussed  in 
a  subsequent  chapter. 

The  fofm  of  apparatus 
used  in  determining  the  re- 
lation between  the  volumes 
of  hydrogen  and  oxygen  unit- 
ing and  that  of  the  aqueous 
vapor  formed  is  illustrated 
in  Fig.  33.  The  arm  A  of 
the  eudiometer  in  which  the 
combination  of  the  gases  is 
effected  is  surrounded  by 
a  tube  through  which  is 
passed  steam  or,  preferably,  • 
the  vapor  of  some  liquid 
boiling  above  100°  (amyl 

alcohol  is  often  used).  A  mixture  of  2  volumes  of  hydrogen  with  1  of  oxygen  is 
introduced  into  the  eudiometer.  A  suitable  liquid  is  then  boiled  in  the  flask  B. 
The  resulting  vapor  is  conducted  through  the  space  between  the  tube  A  and  the 
outer  tube,  and  is  then  condensed  as  shown  in  the  figure.  When  the  volume  of 
the  mixed  gases  in  A  has  become  stationary,  showing  that  the  temperature  of 
the  gases  is  the  same  as  that  of  the  vapor,  and  the  pressure  adjusted  as  in  the 
former  experiment  (see  Fig.  30  and  description),  the  reading  on  the  eudiometer 
tube  is  noted.  The  union  of  the  two  gases  is  then  effected  by  an  electric  spark 
from  an  induction  coil  C,  the  pressure  is  adjusted,  -and  the  reading  again  noted 
after  the  volume  of  the  vapor  has  become  constant.  The  volume  of  the  vapor 
thus  obtained  can  be  compared  directly  with  the  volumes  of  the  hydrogen  and 
oxygen  which  united  to  form  it. 


FIG. 33 


HYDROGEN  PEROXIDE 

Composition.  In  1818,  while  studying  the  action  of  acids  upon 
certain  oxides,  the  French  chemist  Thenard  discovered  the  compound 
which  we  now  call  hydrogen  peroxide,  or  sometimes  hydrogen  dioxide. 
The  pure  compound  is  a  liquid  and,  like  water,  is  composed  of 
hydrogen  and  oxygen.  The  proportions  in  which  the  hydrogen  and 


GENERAL  CHEMISTRY 


oxygen  are  present  in  these  two  compounds,  however,  are  widely 
different,  as  shown  in  the  following  statement: 

Water 1  part  of  hydrogen  to    7.94  parts  of  oxygen  by  weight. 

Hydrogen  peroxide  .     .     1  part  of  hydrogen  to  15.88  parts  of  oxygen  by  weight. 

In  other  words,  the  weight  of  oxygen  combined  with  a  fixed  weight  of 
hydrogen  is  just  twice  as  great  in  hydrogen  peroxide  as  in  water. 
This  larger  percentage  of  oxygen  is  indicated  by  the  name  peroxide, 
the  prefix  per  meaning  "  more  "  or  "  excess." 

Preparation.  While  a  dilute  solution  of  hydrogen  peroxide  may 
be  easily  obtained,  the  pure  compound  cannot  be  prepared  without 
great  difficulty,  since  it  readily  decomposes  into  water  and  oxygen. 
Dilute  solutions  of  the  compound  are  prepared  by  the  action  of  acids 
upon  certain  oxides  in  the  presence  of  water.  The  oxide  commonly 
used  is  barium  peroxide.  For  the  acids,  one  may  conveniently  use 
either  sulfuric  or  phosphoric  acid,  properly  diluted  with  water.  With 
sulf uric  acid  the  reaction  may  be  represented  as  follows  : 

sulfuric  acid  -f-  barium  peroxide  =  barium  sulf  ate  +  hydrogen  peroxide 

[  hydrogen!  ("barium"! 

J.,    l  [barium! 

sulfur  sulfur 

I  oxygen  I 
oxygen    J  [oxygen J 

It  will  be  noted  that  in  this  reaction  the  barium  of  the  barium  peroxide 
changes  places  with  the  hydrogen  of  the  acid.  The  barium  sulfate 

formed  is  insoluble,  while  the  hydrogen  per- 
oxide dissolves  in  the  water  present.  The 
barium  sulfate  may  therefore  be  removed 
from  the  solution  by  filtration.  Phosphoric 
acid  acts  in  a  similar  way.  In  this  way  one 
can  readily  prepare  a  dilute  solution  of  the 
peroxide  in  water.  To  concentrate  this,  the 
solution  is  transferred  to  a  separatory  fun- 
nel (Fig.  34),  ether  is  added,  and  the  con- 
tents thoroughly  shaken.  The  hydrogen 
peroxide,  being  more  soluble  in  ether  than 
in  water,  is  largely  dissolved  by  the  ether. 
On  standing,  the  ether  rises  to  the  surface 
of  the  water,  carrying  with  it  the  dissolved  hydrogen  peroxide.  This 
solution  is  then  separated  from  the  water  by  allowing  the  latter  to  run 
out  through  the  stopcock  A,  and  the  ether  is  evaporated.  Since  the  boiling 


Phydrogen~] 
[oxygen     J 


FIG.  34 


WATER  71 

point  of  ether  is  34.6°,  it  is  possible  to  remove  it  at  a  comparatively  low 
temperature,  and  thus  to  prevent  any  marked  decomposition  of  the 
hydrogen  peroxide.  By  distilling  off  the  ether  under  diminished  pres- 
sure the  separation  may  be  effected  at  a  still  lower  temperature.  In 
this  way,  with  proper  precaution,  one  can  obtain  a  solution  containing 
nearly  90  per  cent  of  the  peroxide.  On  cooling  this  solution  with  a 
freezing  mixture,  clear  crystals  of  the  pure  hydrogen  peroxide  {melting 
at  —  2°)  separate.  As  in  the  case  of  ozone,  however,  the  pure  sub- 
stance is  seldom  prepared,  because  of  its  highly  explosive  character. 

Properties.  Hydrogen  peroxide  is  a  clear,  sirupy  liquid  having  a 
density  of  1.458.  It  cannot  be  distilled  under  ordinary  atmospheric 
pressure,  since  it  decomposes  into  water  and  oxygen  with  explosive 
violence  before  the  boiling  point  is  reached.  Under  greatly  diminished 
pressure,  however,  it  may  be  distilled  with  little  decomposition  ;  thus 
at  a  pressure  of  29  mm.  it  boils  at  69°.  It  mixes  with  water,  ether, 
and  alcohol  in  all  proportions. 

Since  hydrogen  peroxide  so  readily  decomposes,  with  evolution  of 
oxygen,  it  acts  as  a  strong  oxidizing  agent,  even  in  very  dilute  solu- 
tions. An  easily  oxidizable  substance,  like  wool,  is  ignited  by  the 
addition  of  a  few  drops  of  the  pure  compound.  The  speed  of  decom- 
position of  hydrogen  peroxide  is  influenced  in  many  ways.  In  dilute 
solutions  and  at  a  low  temperature  the  speed  is  very  slow,  while  at 
higher  temperatures  and  in  more  concentrated  solutions  it  becomes  so 
great  as  to  cause  violent  explosions.  Moreover,  the  speed  of  decom- 
position is  greatly  affected  by  the  presence  of  certain  catalytic  agents. 
Thus  a  little  finely  divided  platinum  or  manganese  dioxide,  added  to 
a  concentrated  solution  of  the  peroxide,  produces  such  rapid  decom- 
position as  to  cause  an  explosion.  Certain  organic  substances  have 
a  similar  action.  Just  as  some  substances  increase  the  rapidity  of 
decomposition,  so  others  retard  it.  Thus  the  ordinary  solution  of  hy- 
drogen peroxide  sold  for  medicinal  purposes  contains  a  small  amount 
of  some  such  substance,  generally  a  trace  of  acid,  which  is  added  to 
preserve  the  strength  of  the  solution  by  retarding  decomposition. 

The  strong  oxidizing  properties  of  hydrogen  peroxide  may  be  shown  by  its 
action  upon  lead  sulfide.  This  is  a  black  compound  of  lead  and  sulfur.  When 
treated  with  hydrogen  peroxide  it  is  oxidized  to  a  compound  of  lead,  sulfur,  and 
oxygen  known  as  lead  sulf  ate,  which  is  white.  Through  the  action  of  the  hydro- 
gen peroxide,  therefore,  the  black  color  of  the  lead  sulfide  gradually  gives  way  to 
the  white  color  of  the  lead  sulfate. 


72  GENERAL  CHEMISTRY 

Uses.  Hydrogen  peroxide  has  many  commercial  uses,  all  based 
on  its  strong  oxidizing  properties.  The  common  medicinal  peroxide 
of  the  druggist  is  an  aqueous  solution  containing  3  per  cent,  by 
weight,  of  the  compound.  The  efficiency  of  hydrogen  peroxide  as 
a  germicide  is  due  to  the  oxygen  liberated,  which  destroys  any  micro- 
organisms present.  Like  ozone,  it  acts  upon  certain  dyes  and  natural 
colors,  such  as  that  of  the  hair,  oxidizing  them  into  colorless  com- 
pounds ;  hence  it  is  sometimes  used  as  a  bleaching  agent.  The  chemist 
finds  it  especially  useful  as  an  oxidizing  agent  in  many  analytical 
operations.  For  this  purpose  it  is  often  convenient  to  have  a  rather 
concentrated  solution,  so  that  a  30  per  cent  solution  is  now  sold  as 
a  commercial  product. 


CHAPTER  VI 

THE  THREE  STATES  OF  MATTER 

The  states  of  matter.  The  study  of  water  has  brought  to  our  atten- 
tion a  substance  existing  in  three  very  different  states,  namely,  gase- 
ous, liquid,  and  solid.  In  a  general  way  a  gas  or  vapor  may  be  regarded 
as  matter  in  such  a  state  that  it  distributes  itself  uniformly  through- 
out the  space  in  which  it  is  placed.  A  liquid  does  not  so  distribute 
itself;  it  has  no  characteristic  shape  of  its  own,  however,  but  takes 
the  form  of  the  vessel  in  which  it  is  placed.  A  solid  retains  its  own 
shape  irrespective  of  the  size  or  form  of  the  containing  vessel.  The 
occurrence  in  all  of  these  three  states  is  not  peculiar  to  water  but  is 
found  to  be  true  of  the  great  majority  of  substances.  It  will  therefore 
be  of  advantage  to  obtain  a  more  accurate  idea  as  to  the  difference 
between  these  states,  and  of  the  conditions  under  which  a  substance 
may  pass  from  one  to  another. 

1.  The  relation  between  liquids  and  gases.  We  shall  first  consider  the 
relation  between  liquids  and  gases. 

Evaporation.  When  a  liquid  such  as  water  is  placed  in  an  open 
vessel,  it  gradually  passes  into  the  air  in  the  form  of  gas,  the  process 
being  called  evaporation.  If  it  is  in  a  confined  space,  which  it  only 
partially  fills,  as,  for  example,  in  a  closed  bottle,  evaporation  proceeds 
until  the  air  above  the  liquid  contains  a  definite  percentage  of  gaseous 
water,  and  then  apparently  ceases,  the  air  being  said  to  be  saturated 
with  water  vapor.  From  the  kinetic  point  of  view  we  may  imagine 
the  particles  of  the  liquid  to  be  in  motion  but  moving  more  slowly 
than  in  a  gas,  with  very  frequent  collisions  and  subject  to  very  con- 
siderable mutual  attraction.  From  the  surface  of  the  liquid  the  more 
rapidly  moving  particles  will  from  time  to  time  escape,  breaking  free 
from  the  attraction  of  the  liquid.  They  will  then  move  about  in  the 
the  space  above  it  as  gas  particles,  and  will  from  time  to  time  return 
to  the  liquid.  When  the  rate  at  which  they  escape  is  equal  to  the 
rate  at  which  they  return,  an  equilibrium  will  be  reached  and  there 
will  apparently  be  no  further  evaporation. 

73 


74  GENERAL  CHEMISTEY 

Effect  of  temperature  upon  evaporation.  If  the  liquid  is  now  warmed, 
the  rate  of  motion  of  its  particles  is  increased  and  consequently  the 
rate  of  evaporation  is  increased.  A  new  equilibrium  is  reached  at  the 
higher  temperature,  there  being  a  larger  percentage  of  the  gaseous  sub- 
stance in  the  air  than  before.  The  quantity  of  water  present  in  gaseous 
form  over  water  in  a  closed  space  is  therefore  roughly  proportional 
to  the  temperature,  but  experiment  shows  that  it  is  not  accurately  so. 
For  exact  values  we  must  consult  tables  based  on  experiment,  or 
curves  plotted  from  such  experiments. 

Vapor  pressure.  The  quantity  of  water  present  in  air  in  gaseous 
form  may  be  expressed  in  a  number  of  ways.  The  Weather  Bureau 
expresses  it  as  relative  humidity,  meaning  by  this  term  the  quantity 
present  as  compared  with  the  quantity  which  would  be  present  at 
the  same  temperature  when  equilibrium  is  reached.  It  might  be  stated 
simply  as  the  weight  in  grams  in  a  liter  of  air.  A  more  satisfactory 
way  is  to  express  the  quantity  in  terms  of  the  pressure  which  it  exerts 
as  a  gas.  Of  the  total  pressure  of  the  air  on  the  surface  of  the  water 
a  part  is  due  to  oxygen,  a  part  to  nitrogen,  and  a  part  to  gaseous 
water.  We  may  therefore  indicate  the  fraction  of  the  total  pressure 
due  to  gaseous  water,  and  so  have  a  convenient  method  of  expressing 
the  quantity  of  water  present  in  the  air  which  is  independent  of  the 
volume.  Owing  to  this  method  of  expression  the  value  of  the  water 
in  the  gaseous  form  at  a  given  temperature  is  called  the  vapor  pressure, 
or  aqueous  tension,  of  the  water.  The  vapor  pressure  of  a  liquid  may 
therefore  be  defined  as  the  pressure  upon  its  surface  due  to  its  own 
vapor.  A  table  showing  the  vapor  pressure  of  water  at  various  tem- 
peratures will  be  found  in  the  Appendix. 

For  example,  air  confined  over  water  at  20°  will  take  up  water  until  the  water 
vapor  exerts  a  gas  pressure  equal  to  17.4  mm.  If  the  total  pressure  of  the  air  is 
then  760  mm.,  17.4  mm.  is  due  to  water  vapor  and  742.6  mm.  to  other  gases. 

17  4 

Therefore  — ^—  of  the  total  volume  of  the  air  is  water  vapor.    If  this  volume  is 
7oO 

1000  cc.,  — —  x  1000  =  22.189  cc.  is  the  volume  the  water  vapor  would  itself  occupy 

i  uU 

as  a  gas  at  760  mm.  pressure  and  at  20°. 

Correction  for  vapor  pressure  in  gas  measurements.  If  a  gas  is  col- 
lected over  water,  it  is  evident  that  the  observed  volume  includes  not 
only  that  of  the  gas  but  also  that  of  the  water  present  as  vapor.  This 
latter  volume  could  be  calculated  in  the  way  just  indicated,  and 


THE  THREE  STATES   OF  MATTER, 


75 


deducted  from  the  observed  volume  of  the  gas.  A  more  convenient 
method,  especially  if  the  gas  volume  is  to  be  reduced  to  standard 
conditions,  is  to  subtract  the  value  of  the  vapor  pressure  at  the  ob- 
served temperature  from  the  pressure  under  which  the  gas  is  measured. 
The  formula  for  reduction  to  standard  conditions  (p.  49)  will  then  be 

^  (P  -  g)  x  V  X  273  ^ 
760  x  T 

in  which  a  is  the  value  of  the  vapor  pressure  of  water  at  the  temper- 
ature of  observation. 

Determination  of  vapor  pressure.  Experimentally  the  value  of  the  vapor  pres- 
sure of  a  liquid  at  any  temperature  may  be  determined  in  the  following  way : 
Two  long  barometer  tubes  are  filled  with  mercury  and 
inverted  in  an  open  vessel  of  the  same  liquid  (Fig.  35). 
A  few  drops  of  the  liquid  to  be  examined  are  introduced 
under  the  open  end  of  one  of  the  tubes,  the  liquid  so 
introduced  immediately  rising  to  the  top  of  the  mercury 
column.  Evaporation  at  once  takes  place  and,  because 
of  the  pressure  of  the  gas  so  formed,  the  mercury  column 
falls  to  some  extent.  When  equilibrium  is  reached,  the 
difference  in  level  of  the  mercury  in  the  two  tubes, 
included  between  the  dotted  lines  A  and  B  in  the  figure, 
will  correspond  to  the  vapor  pressure  of  the  liquid  ex- 
pressed in  millimeters  of  mercury.  The  tubes  may  be 
surrounded  by  jackets  through  which  heated  liquids 
are  circulated,  so  that  any  desired  temperature  may  be 
secured. 


-A 
B 


FIG.  35 


Boiling  point.  During  the  heating  of  a  liquid  a 
portion  of  the  energy  imparted  to  it  goes  to  raise 
its  temperature  and  a  portion  to  change  it  into 
a  vapor  at  its  surface.  When  the  pressure  of  the  vapor  arising  from 
the  liquid  just  exceeds  the  opposing  atmospheric  pressure,  all  of  the 
heat  energy  goes  to  change  the  liquid  into  vapor  and  into  mechanical 
work  in  pushing  back  the  atmosphere,  and  the  temperature  remains 
constant.  This  temperature  is  called  the  boiling  point  under  the  pres- 
sure in  question.  Since  the  boiling  point  depends  upon  the  atmospheric 
pressure,  it  is  necessary  to  adopt  a  standard  pressure  under  which  it 
shall  be  measured,  and  this  is  taken  as  that  of  a  column  of  mercury 
760  mm.  in  height.  Under  760  mm.  pressure,  water  boils  at  100° ; 
under  a  pressure  of  525.5  mm.,  at  90°. 

We  usually  think  of  a  liquid  as  boiling  when  bubbles  form  and  rise  freely 
through  it,  and  this  is  really  an  accurate  test,  as  will  be  seen  by  reference  to 


76 


GENERAL   CHEMISTRY 


760mm. 


Fig.  36.  The  figure  represents  a  vessel  in  which  water  is  being  heated  under 
an  atmospheric  pressure  of  760  mm.  Suppose  a  bubble  of  water  vapor  is  formed 
at  A.  The  pressure  upon  the  bubble  will  be  760  mm.  plus  the  weight  of  the 

water  above  it.  In  order  that  the  bub- 
ble may  survive,  the  pressure  of  the 
vapor  within  it  upon  the  inclosing 
water  must  exceed  the  pressure  of  the 
water  upon  the  bubble.  At  50°  (the 
side  B)  a  bubble  forming  at  the  bot- 
tom and  moving  up  through  the  colder 
liquid  has  not  enough  vapor  pressure 


50°  100' 

FIG.  36 

to  balance  the  opposing  pressure,  and 

it  gradually  collapses.  At  100°  (side  C)  the  vapor  pressure  of  the  bubble  exceeds 
the  external  pressure  ;  it  is  not  cooled  as  it  rises,  but  increases  in  size  as  the  pres- 
sure diminishes^  and  ffnally  escapes  from  the  surface.  The  formation  and  escape 
of  vapor  from  within  the  liquid  absorbs  all  of  the  heat  applied,  and  there  can 
be  no  further  rise  of  temperature  at  that  pressure. 

Heat  of  vaporization  and  condensation.  The  quantity  of  heat  absorbed 
in  changing  1  g.  of  a  liquid  at  its  boiling  point  into  1  g.  of  vapor  at 
the  same  temperature  is  called  the  heat  of  vaporization.  For  water 
this  is  unusually  large  and  amounts  to  539  cal.  Conversely,  if  a  gas 
is  maintained  at  a  pressure  of  760  mm.  and  is  gradually  cooled,  con- 
densation to  the  liquid  state  begins  when  the  boiling  point  is  reached. 
During  liquefaction  the  temperature  remains  constant,  and  a  quantity 
of  heat  is  given  out  exactly  equal  to  the  heat  of  vaporization.  This  is 
called  the  heat  of  condensation. 

Critical  point.  At  high  temperatures  gases  may  be  subjected  to  the 
greatest  possible  pressure  without  passing  into  the  liquid  state.  Thus 
if  steam  is  heated  above  365°,  thousands  of  atmospheres  of  pressure 
will  not  liquefy  it.  If,  while  holding  the  gas  under  a  high  pressure, 
the  temperature  is  slowly  lowered,  a  definite  point  is  reached  at  which 
the  gas  suddenly  liquefies.  This  temperature  is  called  the  critical 
temperature,  and  the  pressure  required  to  cause  liquefaction  at  the 
critical  temperature  is  called  the  critical  pressure. 


BOILING  POINT 

CRITICAL 
TEMPERATURE 

CRITICAL  PRESSURE 

Hvdrosren 

—  252.7° 

—  234.5° 

20  00  atmospheres 

Nitrogen                      ... 

—  195.7° 

—  146° 

33  00  atmospheres 

Oxvcren 

—  182.9° 

—  119° 

50  00  atmospheres 

Carbon  dioxide     .... 
Water 

-    79.0° 
+  100  0° 

+    31.35° 
+  365° 

72.90  atmospheres 

THE  THEEE   STATES  OF  MATTER 


77 


The  critical  pressure  is  not  in  general  more  than  about  60  atmos- 
pheres. If  100  atmospheres  pressure  does,  not  liquefy  a  gas,  it  is 
probably  useless  to  increase  the  pressure,  a  lower  temperature  being 
the  necessary  condition.  The  relations  of  critical  temperature  and 
pressure  were  discovered  by  Andrews  in  1869,  during  his .  researches 
on  carbon  dioxide.  Prior  to  that  time  thousands  of  atmospheres  of 
pressure  had  been  applied  to  oxygen  gas  in  an  effort  to  liquefy  it, 
but  to  no  purpose,  since  it  was  above  its  critical  temperature. 

Methods  of  liquefaction  of  gases.  The  earliest  systematic  efforts  at 
liquefaction  of  gases  were  those  of  the  English  scientist  Faraday,  be- 
ginning about  1823.  He  relied  upon  the  effect  of  pressure  together 
with  moderate  cooling,  most  of 
his  experiments  being  carried  out 
in  the  following  way :  A  quantity 
of  solid  material  which,  when 
heated,  would  liberate  a  consid- 
erable quantity  of  the  gas  to  be 
liquefied,  was  placed  in  one  end 
of  a  bent  tube.  The  other  end 
was  sealed  and  the  tube  arranged 
as  shown  in  Fig.  37,  A  being  the 
solid  material  and  B  a  bath  of  ice 
water.  Upon  heating,  the  gas  is 

given  off  in  a  confined  space  and,  being  under  great  pressure,  lique- 
fies in  the  cold  portion  of  the  tube.  In  this  general  way  Faraday 
liquefied  a  number  of  gases  such  as  ammonia  and  carbon  dioxide. 

Later  experimenters  made  use  of  much  lower  temperatures.  These 
low  temperatures  were  secured  by  taking  advantage  of  the  heat  of 
vaporization  of  low-boiling  liquids.  When  a  quantity  of  any  gas,  such 
as  sulfur  dioxide,  is  liquefied  by  pressure,  cooled,  and  then  allowed 
to  boil  away  under  the  pressure  of  the  atmosphere  or  in  a  partial 
vacuum,  the  temperature  of  the  liquid  falls  to  its  boiling  point  under 
that  pressure.  The  very  cold,  boiling  liquid  may  be  used  as  a  bath 
to  cool  some  other  gas  below  its  critical  temperature,  when  it  may 
in  turn  be  liquefied  by  pressure.  By  employing  such  a  process, 
Cailletet,  in  1877,  first  liquefied  oxygen. 

Since  1895  purely  mechanical  methods  have  been  employed  in 
liquefying  such  gases  as  air.  Machines  constructed  for  this  purpose 
depend  for  their  efficiency  upon  the  cooling  effect  produced  when  a 


FIG. 37 


78 


GENERAL  CHEMISTRY 


From  Compressor 


A- 


FIG.  38 


highly  compressed  gas  is  allowed  to  expand  freely.  When  a  gas  is 
compressed,  heat  is  liberated,  and  when  it  is  allowed  to  expand,  heat 
is  absorbed.  It  is  found  that  the  heat  absorbed  is  slightly  greater 
than  that  liberated,  owing  to  the  fact  that  gases  do  not  exactly  con- 
form to  the  gas  laws.  When  a  gas  is  alternately  compressed  and 

expanded,  a  cooling  bath  being  maintained 
around  it  to  take  up  the  heat  of  compres- 
sion, its  temperature  steadily  falls  to  the 
point  of  liquefaction. 

The  Linde  machine.  In  the  Linde  machine 
(Fig.  38)  the  compression  is  effected  by  a  strong 
pump.  The  compressed  air  at  200  atmospheres 
pressure  is  first  cooled  in  a  freezing  bath  A.  It 
then  passes  upward  as  indicated  by  the  arrow  and 
enters  the  inner  tube  of  a  system  of  three  concen- 
tric, spirally  wound  copper  tubes.  At  the  lower 
end  of  this  system  it  expands  through  a  valve 
operated  by  the  head  screw  E  to  a  pressure  of 
from  20  to  50  atmospheres,  and  in  so  doing  be- 
comes much  colder.  It  is  then  returned  to  the 
pump  through  the  space  between  the  inner  and 

second  tube  and  the  pipe  at  the  top,  cooling  the  interior,  compressed  gas.  When 
this  process  no  longer  results  in  a  fall  of  temperature,  the  valve  C  is  opened, 
whereby  some  of  the  cold  air  at  20  atmospheres  pressure  is  allowed  to  expand  to 
atmospheric  pressure.  In  so  doing  a  part  liquefies  and  is  caught  in  the  vessel  D, 
while  the  very  cold  air  which  escapes  liquefaction  is  led  back  through  the  outer 
tube  of  the  spiral  to  further  cool  the  air  within  the  two  inner  tubes. 

Dewar  flasks.  For  collecting  and  temporarily  preserving  such 
liquids  Dewar  employed  a  special  type  of  vessel  which  has  come 
to  be  known  by  his  name.  This  consists  of  two  concen- 
tric vessels  of  any  convenient  shape,  such  as  the  one 
shown  in  Fig.  39.  The  two  flasks  are  joined  together  at 
the  upper  rim  only,  and  the  space  between  them  is  ex- 
hausted by  an  air  pump.  A  vacuum  serves  as  the  best 
possible  insulator  to  heat  conduction,  and  the  surface 
of  the  outer  flask  may  also  be  silvered,  so  that  external 
heat  may  be  reflected  from  it  and  not  absorbed.  Liquid 
air  may  be  preserved  in  such  a  vessel  for  many  hours. 
Vessels  of  the  same  plan  of  construction  are  now  sold  for  the  purpose 
of  keeping  liquids  either  hot  or  cold  during  long  journeys,  and  are 
very  effective. 


FIG.  39 


THE   THEEE   STATES   OF  MATTER  79 

2.  The  relation  between  liquids  and  solids.  Let  us  now  consider 
the  relation  between  liquids  and  solids. 

Solid  bodies.  When  the  majority  of  liquids  are  cooled  sufficiently, 
they  reach  a  temperature  at  which  a  sudden  change  begins,  the  liquid 
gradually  freezing  to  a  mass  of  crystals,  while  the  temperature  remains 
constant.  With  water  this  takes  place  at  zero.  Some  liquids,  such  as 
waxes,  glasses,  and  glues,  simply  become  less  and  less  fluid  and  do 
not  have  any  definite  point  of  solidification.  They  finally  reach  a  con- 
dition in  which  they  are  apparently  solids.  It  has  been  found  that 
those  solid  substances  which  have  a  crystalline  structure  have  a  sharp 
solidifying  point,  while  noncrystalline,  or  amorphous,  substances  pass 
gradually  from  undoubted  liquids  into  what  appear  to  be  rigid  solids, 
with  no  sharp  point  of  transition.  These  are  best  regarded  as  still 
liquids  but  so  viscous  as  to  be  quite  rigid,  the  term  solid  being 
reserved  for  crystalline  bodies  having  a  sharp  solidifying  point. 

Melting  point.  When  a  crystalline  solid  is  slowly  heated,  the  tem- 
perature steadily  rises  to  a  certain  definite  point.  Further  application 
of  heat  does  not  raise  the  temperature,  but  the  solid  begins  to  melt. 
The  temperature  remains  constant  until  the  melting  is  complete,  and 
then  rises  again.  The  heat  energy  supplied  all  through  the  melting 
has  no  effect  on  the  temperature,  but  is  used  up  in  altering  the  physi- 
cal state  of  the  substance.  The  quantity  of  heat  absorbed  in  convert- 
ing 1  g.  of  a  solid  at  its  melting  point  into  1  g.  of  liquid  at  the 
same  temperature  is  called  the  heat  of  fusion  of  the  substance.  For 
water  this  amounts  to  79  cal. 

Freezing  point.  When  a  liquid  is  cooled  it  does  not  always  begin 
to  solidify  when  the  melting  point  is  reached.  It  may  be  cooled  con- 
siderably below  that  point,  and  the  liquid  is  then  said  to  be  undercooled. 
This  sometimes  happens  to  water  in  a  shallow  pool  on  a  cold,  still 
night.  If  now  a  fragment  of  the  solid  is  placed  in  the  liquid,  or  if  the 
liquid  is  violently  shaken,  solidification  at  once  begins,  the  temperature 
rising  to  the  true  freezing  point  and  remaining  constant.  The  freezing 
point  is  therefore  most  accurately  defined  as  the  temperature  at  which 
the  liquid  and  the  solid  will  remain  unchanged  in  contact  with  each 
other.  The  more  viscous  a  liquid  is  at  its  freezing  point  the  more 
readily  undercooling  takes  place ;  and  with  very  viscous  liquids,  as  we 
have  seen,  true  solidification  may  never  occur. 

In  the  process  of  solidification  heat  is  given  out  corresponding  to 
the  heat  of  fusion.  In  both  cases  it  will  be  noticed  that  the  energy 


80  GENERAL  CHEMISTRY 

change  opposes  the  physical  change  taking  place.  The  evolution  of 
heat  during  solidification  retards  the  freezing,  for  it  is  only  as  this 
heat  is  lost  by  radiation  that  the  solidification  can  continue.  Were 
it  not  for  this,  ponds  would  freeze  very  rapidly  in  winter  when  the 
freezing  point  is  reached.  Amorphous  bodies  have  no  point  at  which 
we  can  detect  a  corresponding  evolution  or  absorption  of  heat, 
which  is  another  reason  for  regarding  them  as  liquids,  even  when 
they  are  apparently  solids. 

Vapor  pressure  of  solids.  Many  solids  give  off  vapor  at  ordinary 
temperatures,  just  as  do  liquids.  This  is  evident  from  the  odor  of 
such  solids  as  camphor  and  naphthalene  (moth  balls).  As  the  solid  is 
heated  this  vapor  pressure  increases  in  value.  If  it  increases  to  the 
point  where  it  just  exceeds  the  pressure  of  the  atmosphere,  the  solid 
passes  directly  into  a  gas  without  melting  or  boiling.  This  is  the 
case  with  quite  a  number  of  solids,  such  as  arsenic  and  ammonium 
chloride  (sal  ammoniac).  When  the  vapors  from  such  solids  are 
cooled,  they  pass  directly  back  into  the  solid  form.  The  process  of  con- 
verting a  solid  into  a  vapor  and  cooling  the  vapor  to  a  solid  again  is 
called  sublimation,  and  the  solid  is  said  to  sublimate  on  heating.  It 
will  be  remembered  that  the  corresponding  process  with  liquids  is 
called  distillation.  Solids  which  have  a  sufficient  vapor  pressure  are 
often  separated  from  nonvolatile  impurities  by  sublimation,  and  this 
has  given  rise  to  such  names  as  corrosive  sublimate  (mercuric  chloride). 
Transformation  diagram.  The  several  transformations  of  a  substance 
can  be  represented  most  conveniently  in  the  form  of  a  diagram  in 
which  temperature  and  pressure  are  taken  as  the  coordinates.  Such  a 

diagram  is  shown  in  Fig.  40.  The  curve 
OA  represents  the  increase  of  vapor  pres- 
sure of  the  liquid  with  rise  in  temperature, 
the  vapor  pressures  pvp2  corresponding  to 
the  temperatures  tv  t2.  This  curve  ends 
at  a  point  A,  which  is  the  critical  temper- 
ature at  which  the  distinction  between 


FIG  4Q  liquid  and  vapor  abruptly  ceases.    OB  is 

the   vapor-pressure  t  curve    of   the    solid, 

usually  called  the  sublimation  curve,  the  pressure  p  corresponding  to 
the  temperature  t.  The  point  0,  where  these  two  curves  intersect,  is 
the  freezing  point  of  the  liquid,  the  vapor  pressures  of  solid  and 
liquid  being  equal.  Since  the  two  curves  intersect  at  this  point,  the 


THE  THREE  STATES  OF  MATTER 


81 


solid  and  liquid  can  coexist  at  -  this  temperature  and  pressure,  but 
they  can  do  so  at  no  other  point;  this  is  therefore  called  a  transi- 
tion point.  If  no  solid  makes  its  appearance  at  the  freezing  point,  the 
vapor  pressure  of  the  liquid  will  be  represented  by  the  extension  of 
the  curve  OA  toward  C.  It  will  be  seen  that  the  vapor  pressure  of 
undercooled  liquid  is  greater  than  that  of  the  solid  at  the  same  tem- 
perature, as  at  £,  indicating  a  more  unstable  condition.  Very  few 
solids  can  be  heated  above  the  melting  point  without  melting,  so 
the  curve  OB  can  rarefy  be  prolonged  beyond  the  transition  point. 

Crystals.  When  a  liquid  freezes,  it  changes  into  a  mass  of  solid 
bodies,  each  of  which  has  a  definite  geometric  form  and  is  known  as  a 
crystal.  Similar  bodies  may  also  be  deposited  from  solutions,  or  be 
formed  by  condensing  vapors.  Crystals  are  always  bounded  by  plane 
surfaces,  which  are  arranged  in  an  orderly  fashion  with  reference  to 
imaginary  lines  drawn  through  the  crystal  and  called  its  axes.  Every 
crystal  has  therefore  a  definite  geometric  form.  While  the  variety  of 
form  which  crystals  may  assume  is  almost  endless,  it  has  been  found 
that  they  may  all  be  referred  to  one  of  six  fundamental  arrangements 
of  axes,  these  constituting  what  are  known  as  the  systems  of  crystal- 
lography. These  arrangements,  together  with  two  of  the  simplest 
crystal  forms  referred  to  each,  are  shown  in  the  accompanying  figures 
(Figs.  41-47). 

The  crystal  systems.  The  relation  of  the  axes  in  the  several  systems  is  as 
follows : 

1.  Isometric  or  regular  system  (Fig.  41)  :  three  equal  axes  all  at  right  angles 
to  each  other. 

2.  Tetragonal  system  (Fig.  42)  :  two  equal  axes  and  a  third  of  different  length, 
all  at  right  angles. 

3.  Orthorhombic  system  (Fig.  43)  :  three  unequal  axes  all  at  right  angles. 


FIG.  41 


FIG.  42 


FIG.  43 


4.  Monoclinic  system  (Fig.  44)  :  two  axes  at  right  angles  and  a  third  at  right 
angles  to  one  of  these  but  inclined  toward  the  other.    The  axes  may  be  of  any 
relative  lengths,  and  the  angle  of  inclination  may  vary  from  0°  to  90°. 

5.  Triclinic  system  (Fig.  45)  :  three  axes,  all  inclined  toward  each  other.   The 
axes  may  be  of  any  relative  length,  and  the  angles  of  inclination  may  also  vary. 


82 


GENERAL  CHEMISTRY 


6.  Hexagonal  system  (Fig.  46)  :  three  equal  axes  in  the  same  plane,  intettsect- 
ing  at  angles  of  60°,  and  a  fourth  at  right  angles  to  all  of  these.  In  addition  to 
the  two  general  forms  shown  in  Fig.  46  there  are  many  rhombic  forms  belonging 
to  this  system,  such  as  the  one  represented  in  Fig.  47. 


FIG.  44 


FIG. 45 


FIG.  46 


FIG. 47 


Structure  of  crystals.  There  is  little  doubt  that  these  plans  of  formation 
correspond  to  orderly  arrangements  of  the  particles  of  solid  matter  of  which  the 
crystals  are  composed,  so  that  the  crystals  resemble  in  structure  the  piles  of  can- 
non balls  in  a  military  park.  In  accordance  with  this  idea  it  is  known  that  the 
various  properties  of  the  crystal,  such  as  hardness,  strength,  optical  refraction, 
and  conductivity  toward  heat  and  electricity,  differ  in  different  directions  through 
the  crystal.  Crystals  also  split  in  definite  directions,  giving  plane  surfaces.  Some 
of  these  groupings  represent  a  more  stable  arrangement  than  do  others,  so  that 
when  a  given  substance  crystallizes  in  two  forms,  as  sometimes  happens,  the 
change  from  the  one  to  the  other  is  in  general  accompanied  by  an  energy  change. 
It  is  evident  that  a  body  like  glass  might  be  cut  and  polished  so  as  to  be  an 
exact  copy  of  a  crystal,  but  would  really  not  be  one  at  all,  since  it  would  have 
none  of  the  structure  of  a  crystal. 

Crystal  form  a  characteristic  of  a  substance.  In  general,  under  the 
same  conditions,  a  given  substance  will  always  crystallize  in  a  form 
which  may  be  referred  to  the  same  system  and  with  the  same  ratio  of 
axis  lengths  and  degree  of  inclination.  The  actual  crystal  form  may 
be  quite  different,  however.  For  example,  the  form  may  be  either  a 
cube  or  an  octohedron,  both  of  which  are  referred  to  the  same  axes. 
Not  infrequently  a  substance  may,  under  different  conditions,  assume 
two  forms  in  entirely  different  systems,  and  it  is  then  said  to  be 
dimorphous.  For  example,  one  form  may  occur  when  the  substance 
freezes,  and  another  when  it  is  deposited  from  solution.  Trimorphous 
substances  are  also  known.  When  two  substances  crystallize  in  the 
same  form  and  have  the  same  inclination  of  axes  and  the  same  ratios 
in  their  lengths,  they  are  said  to  be  isomorphous. 


CHAPTER  VII 

THE  LAWS  OF  CHEMICAL  COMBINATION :  THE  ATOMIC  THEORY 

I.  Fundamental  laws  of  themical  combination.  Having  considered 
two  typical  elements,  and  having  gained  some  insight  into  chemical 
reactions  through  a  study  of  the  preparation  of  these  elements  and 
the  combinations  which  they  form  with  each  other  and  with  other 
elements,  we  may  now  go  on  one  step  farther.  What  generalizations 
have  been  reached  in  regard  to  the  characteristics  of  chemical  action  ? 
What  theoretical  ideas  have  been  developed  as  to  the  mechanics  of 
this  action?  These  are  the  questions  which  suggest  themselves  and 
which  we  shall  now  consider.  Aside  from  the  question  of  energy 
relations,  our  purely  material  knowledge  of  chemical  action  may  be 
stated  in  the  form  of  four  general  laws. 

1.  The  law  of  conservation  of  mass.  In  the  earlier  stages  of  the  devel- 
opment of  chemistry  little  importance  was  attached  to  the  relations  by 
weight  between  reacting  substances.  In  a  general  way  it  was  assumed 
that  the  total  weight  remained  constant,  but  as  heat,  light,  and  phlo- 
giston (the  principle  of  combustion)  were  all  considered  to  be  mate- 
rial, and  to  escape  during  action,  apparent  loss  of  weight  was  to  be 
expected. 

Lavoisier  first  clearly  stated  the  principle  of  conservation  of 
mass  in  1785,  attributing  apparent  changes  to  experimental  error. 
Since  his  time  scientists  have  been  accustomed  to  regard  the  law 
as  a  sort  of  axiom,  and  few  experimental  researches  have  been 
undertaken  with  the  express  purpose  of  testing  it,  but  experiments 
carried  out  for  other  purposes  can  be  cited  in  its  support.  Thus  the 
work  of  the  Belgian  chemist  Stas  (1865)  shows  that  in  certain  reac- 
tions the  loss  or  gain  could  not  have  been  more  than  from  2  to  4 
parts  in  100,000. 

In  1906  Landolt  published  the  results  of  a  series  of  experiments 
carried  out  at  Berlin  in  critical  test  of  the  law.  His  general  plan  was 
to  -place  the  materials  which  were  to  act  on  each  other  (generally  in 
solution)  in  the  two  limbs  of  a  glass  vessel  of  the  form  represented  in 

83 


84 


GENERAL   CHEMISTRY 


FIG.  48 


Fig.  48.  The  open  ends  were  then  sealed  off  and  the  vessel  weighed. 
The  vessel  was  inverted  and  the  materials  thus  brought  into  contact 
with  each  other,  and  after  the  reaction  the  vessel  was  again  weighed. 
A  large  number  of  such  experiments  were  carried  out  with  every  refine- 
ment of  skill  and  apparatus,  and  very  slight  differences  between  the 
two  weights  were  detected.  These  were  never  more  than  a  few  hun- 
dredths  of  a  milligram  in  a  total  weight  of  100  g.,  that 
is,  about  1  part  in  10,000,000.  It  is  questionable 
whether  these  slight  differences  exceed  the  unavoid- 
able experimental  error.  Certainly  we  may  state  the 
law  in  the  following  form:  Within  the  limits  of  ex- 
perimental accuracy  no  change  in  the  total 'weight 
of  matter  can  be  detected  when  chemical  action  takes 
place. 

2.  The  law  of  definite  composition.  The  common  ex- 
periences of  the  earlier  chemists  led  them  to  believe 
that  the  composition  of  a  pure  compound  is  quite 
definite.  The  question  as  to  whether  this  is  so  or 
not  became  an  important  issue  in  the  years  1802-1808,  as  a  result  of 
the  views  of  a  distinguished  Frenchman,  Berthollet.  On  theoretical 
grounds  Berthollet  was  led  to  believe  that  the  composition  of  a  sub- 
stance is  somewhat  variable,  being  dependent  on  the  relative  quanti- 
ties of  the  several  materials  present  at  the  time  of  its  formation.  For 
instance,  experiment  showed  that  the  composition  of  iron  sulfide  is, 
at  least  approximately,  iron  63.55  per  cent,  sulfur  36.45  per  cent. 
Berthollet  thought  that  such  figures  were  only  approximate  —  that  if 
equal  parts  of  iron  and  sulfur  were  to  be  heated  together,  a  larger 
percentage  of  sulfur  would  be  present  in  the  product. 

These  views  were  strongly  opposed  by  a  fellow  countryman,  Proust, 
who  was  professor  of  chemistry  at  Madrid  during  most  of  the  con- 
troversy. Proust  maintained  that  the  composition  of  a  pure  compound 
is  perfectly  definite,  and  that  when  two  elements  form  more  than  one 
compound,  each  has  its  own  exact  composition,  there  being  no  inter- 
mediate gradations.  He  maintained  that  apparent  variability  is  due 
to  lack  of  purity  in  the  compound.  Proust's  experimental  work  was 
very  accurate  for  his  time,  but  his  analyses  were  subject  to  errors  of 
from  1  to  2  per  cent.  The  advance  in  experimental  exactness  has 
steadily  demonstrated  the  correctness  of  Proust's  conclusions.  In 
1860  and  again  in  1866  the  Belgian  chemist  Stas  undertook  elaborate 


THE  LAWS  OF  CHEMICAL  COMBINATION  85 

researches  in  a  critical  study  of  the  law  of  constant  composition, 
his  analyses  being  trustworthy  in  some  instances  to  within  about 
1  part  in  50,000.  Within  these  limits  he  showed  that  the  law  holds 
rigidly..  In  our  own  time  the  work  of  the  American  chemist  Theodore 
Richards,  in  a  connection  to  be  mentioned  a  little  later,  has  demon- 
strated the  accuracy  of  the  law  within  still  narrower  limits.  Within 
the  limits  of  modern  methods  of  analysis,  then,  we  may  state  the  law  : 
The  composition  of  a  pure  compound  is  always  precisely  the  same. 

3.  The  law  of  multiple  proportion.  Proust  investigated  the  composi- 
tion of  several  pairs  of  compounds  formed  from  the  same  two  ele- 
ments, and  the  following  table  illustrates  his  results: 


OXIDE  OF  TIN 

OXIDE  OF  COPPER 

Tin 

Oxygen 

Copper 

Oxygen 

78.4% 
87% 

21.6% 
13% 

80% 
86.2% 

20% 
13.8% 

Proust  made  no  comment  on  the  relation  between  the  ratios  of  oxy- 
gen to  metal  in  the  two  cases,  and  his  figures  suggest  none.  Three- 
investigators, —  Dalton,  an  English  school-teacher  (1805),  Wollaston, 
his  fellow  countryman  (1808),  and  Berzelius,  a  Swede  (1811), —  quite 
independently  of  each  other,  observed  a  striking  relationship  in 
such  cases,  which  has  come  to  be  known  as  Dalton's  law  of  multiple 
proportion,  since  Dalton  first  formulated  it  and  was  very  active  in 
seeking  proofs  of  its  validity.  He  showed  that  if  the  composition  in 
such  cases  is  stated  not  in  percentages  but  in  the  weights  of  one  ele- 
ment combined  with  a  fixed  weight  of  the  other,  then  these  weights 
are  in  the  ratio  of  integer  numbers.  He  showed  that  in  nitrous  oxide 
1  part  of  oxygen  is  combined  with  1.648  parts  of  nitrogen,  while 
in  nitric  oxide  1  part  of  oxygen  is  combined  with  0.798  part  of 
nitrogen.  The  ratio  of  the  two  weights  of  nitrogen  is  therefore 
1.648 :  0.798  or  2.06  :  1,  that  is,  2  : 1  within  the  limits  of  error. 

In  the  case  of  the  two  hydrides  of  carbon  (marsh  gas  and  ethylene) 
Dalton  found  that  the  ratios  of  carbon  to  hydrogen  are  respectively 
4.3  carbon :  2  hydrogen  and  4.3  carbon :  1  hydrogen.  He  also  recal- 
culated some  of  Proust's  results,  showing  that  they  supported  his 
generalization,  though  the  deviations  are  as  much  as  5  per  cent. 

The  results  obtained  by  Berzelius  in  quite  a  large  number  of  cases 
showed  the  generalization  to  be  true  to  within  possibly  0.3  per  cent. 


86  GENERAL  CHEMISTRY 

In  more  recent  times  no  definite  tests  of  the  law  have  been  under- 
taken, but  analyses  of  compounds,  made  with  great  care  for  other 
purposes,  have  been  recalculated  to  test  its  accuracy,  and  within  the 
unavoidable  errors  of  analysis  it  has  been  found  to  be  a  precise  state- 
ment of  the  facts.  The  composition  of  the  two  compounds,  water 
and  hydrogen  dioxide,  affords  a  good  illustration  of  this  law  (p.  70), 
which  may  be  conveniently  stated  in  the  following  way  :  When  two 
elements  A  and  B  form  more  than  one  compound,  the  weights  of 
the  element  -4,  which  combine  with  a  fixed  weight  of  the  element  B, 
stand  in  the  ratio  of  small  integers  to  each  other.  The  most  usual 
ratios  are  1  :  1,  1  :  2,  1  :  3,  2  :  3,  and  2  :  5. 

4.  The  law  of  combining  weights.  This  law,  which  is  often  called  the 
law  of  reciprocal  proportion,  was  formulated  by  the  German  chemist 
Richter  as  the  outcome  of  his  researches  between  the  years  1792  and 
1799.  He  was  of  a  mathematical  turn  of  mind,  and  was  interested  in 
studying  the  numerical  relations  between  the  weights  of  combining 
substances.  Most  of  his  studies  were  concerned  with  those  classes 
of  substances  known  as  acids  and  bases,  which  act  readily  upon 
each  other. 

Richter  found  that  if  we  take  a  series  of  acids,  which  we  may  desig- 
nate as  A,  B,  (7,  .  .  .  ,  and  allow  them  to  act  in  succession  upon  a 
series  of  bases,  designated  by  X,  F,  Z,  .  .  .  ,  then  a  simple  relation 
may  be  discovered.  Let  1  g.  of  A,  of  B,  and  of  C  act  successively 
.on  X,  F,  and  Z.  Then 

1  g.  of  A  f  a^  gram  X,       1  g.  of  B  f  ^  gram  X,      1  g.  of  C  (  c^  gram  X, 

unites    \  a2  gram  F,         unites    -j  b2  gram  F,        unites     <{  c2  gram  F, 

with      [  aB  gram  Z  ;          with      [  b3  gram  Z  ;        with       [  cs  gram  Z. 

Then  a^  «2,  a3  are  the  weights  of  X,  F,  Z  respectively,  which  combine 
with  1  g.  of  A  ;  b^  62,  £>8,  the  weights  of  the  same  bases  which  unite 
with  1  g.  of  B,  etc.  Now  Richter  found  that  the  ratio 


That  is,  the  ratio  between  the  several  weights  of  the  bases  X,  F,  Z, 
which  combine  with  a  fixed  weight  of  the  acid  A,  is  the  same  as  the 
ratio  in  which  these  three  bases  combine  with  any  other  acid,  j5,  (7,  D. 
If  with  the  three  bases  X,  F,  Z  this  ratio  had  been  determined  as 
•2.2  :  4.3  :  6.8,  then  this  generalization  of  Richter's  states  that  the 
ratio  in  which  the  three  combine  with  any  acid  is  2.2  :  4.3  :  6.8.  In 


THE  LAWS  OF  CHEMICAL  COMBINATION 


87 


other  words,  2.2  g.  of  the  base  X  has  the  same  value  in  its  action  with 
acids  as  does  4.3  g.  of  Y  or  6.8  g.  of  Z.  If  it  should  be  demonstrated 
that  1  g.  of  an  acid  D  combines  with  5.3  g.  of  the  base  X,  then  we 
can  calculate  what  weight  of  the  base  Y  it  will  require  from  the 
proportion  2.2  :  4.3  : :  5.3  :  x. 

Richter's  work  did  not  have  any  considerable  influence  upon  his  contempo- 
raries, owing  to  a  number  of  causes.  His  language  was  obscure  and  his  ideas 
were  expressed  partly  in  terms  of  the  old  phlogiston  conceptions  and  partly  in 
accord  with  the  newer  oxygen  ideas,  and  so  found  favor  with  the  adherents  of 
neither  theory.  He  was  led  away  from  the  really  important  part  of  his  work  by 
an  endeavor  to  show  that  the  ratios  between  the  combining  numbers  of  the 
bases  are  in  arithmetical  progression,  while  those  of  the  acids  are  in  geometrical 
—  which  is  not  true  at  all. 

Richter's  ideas  were  rediscovered  and  extended  by  other  workers, 
notably  by  Berzelius  in  1811.  It  was  found  that  not  only  to  acids 
and  bases  but  to  every  substance  a  number  can  be  assigned  which 
indicates  its  relative  value  by  weight  in  chemical  reactions.  Evidently 
it  would  be  a  very  great  task  to  determine  by  direct  experiment 
the  combining  number  of  each  known  substance,  but  by  applying 
Richter's  ideas  to  the  elements  the  matter  was  very  greatly  simplified. 

The  combining  weights  of  the  elements.  Experiment  showed  that  it 
is  possible  to  assign  to  each  element  a  number  which  is  proportional  to 
the  weight  by  which  it  enters  into 
chemical  action.  The  meaning  of 
this  statement  is  more  readily  un- 
derstood by  reference  to  the  dia- 
gram (Fig.  49),  which  gives  the 
symbols  of  six  elements,  together 
with  their  combining  weights  as  de- 
termined by  experiment.  By  fol- 
lowing the  line  connecting  any  two 
of  these  elements  we  may  see  at  a 
glance  the  ratio  by  weight  in  which 
they  combine.  Thus  107.88  g.  of 
silver  combines  with  35.46  g.  of 
chlorine,  with  79.92  g.  of  bromine, 
and  with  126.92  g.  of  iodine.  Similarly  100.3  g.  of  mercury  com- 
bines with  79.92  g.  of  bromine,  and  35.46  g.  of  chlorine  combines  with 
126.92  g.  of  iodine.  Sometimes  an  element  acts  upon  a  compound  in 


FIG.  49 


88  GENERAL  CHEMISTRY 

such  a  way  as  to  displace  one  of  the  elements  of  the  compound.  Thus 
copper  will  displace  the  silver  from  silver  nitrate,  and  this  takes  place 
in  such  a  way  that  31.78  g.  of  copper  displaces  107.88  g.  of  silver. 

Elements  with  more  than  one  combining  weight.  The  facts  stated  in 
connection  with  the  law  of  multiple  proportion  at  once  suggest  that 
two  elements  may  combine  in  more  than  one  ratio,  forming  several 
distinct  compounds,  as  in  the  case  of  water  and  hydrogen  peroxide. 
But  since  the  weights  of  the  one  element,  combined  with  a  fixed 
weight  of  the  other,  are  always  in  a  simple  integer  ratio  to  each  other, 
it  follows  that  an  integer  multiple  of  the  combining  weight  of  the 
element  will  always  express  its  combining  value.  Thus  chlorine  forms 
one  compound  with  copper  in  which  35.46  g.  of  chlorine  is  combined 
with  31.78  g.  of  copper,  and  a  second  one  in  which  35.46  g.  of 
chlorine  is  combined  with  63.57  g.  of  copper  (31.78  x  2).  Iodine 
combines  with  chlorine  in  the  ratios  126.92 :  35.46  and  126.92  : 106.38 
(=  35.46  x  3).  We  may  therefore  state  the  law  of  combining  weights 
as  applied  to  elements  thus:  To  each  element  may  be  assigned  a 
number  which  in  itself,  or  when  multiplied  by  some  integer,  repre- 
sents the  weight  by  which  the  element  combines  with  other  elements. 

The  combining  weight  of  a  compound.  Since  the  combining  weights 
of  the  elements  never  change,  save  by  an  integer  multiple,  it  follows 
that  the  combining  weight  of  a  compound  must  be  the  sum  of  the 
combining  weights  of  the  elements  composing  it,  or  of  integer  multi- 
ples of  one  or  more  of  them.  The  combining  weights  of  copper  and 
chlorine  being  31.78  and  35.46  respectively,  the  weight  of  one  of  the 
compounds  is  the  sum  of  these,  namely,  67.24.  The  combining  weights 
of  hydrogen  and  oxygen  are  found  to  be  1.01  and  8.00  respectively, 
water  being  9.01.  When  these  two  compounds  combine,  they  do  so 
in  the  ratio  67.24  :  18.02,  that  is,  once  the  combining  weight  of 
copper  chloride  to  twice  that  of  water. 

II.  The  determination  of  the  combining  weights  of  the  elements. 
Since  the  combining  weight  of  an  element,  or  the  equivalent  weight, 
as  it  is  often  called,  is  thus  found  to  be  a  true  constant  of  nature, 
it  is  of  the  utmost  importance  that  each  one  should  be  determined 
with  the  greatest  possible  precision. 

1.  The  basis  for  the  determination  of  combining  weights.  The  problem 
is  apparently  a  simple  one.  We  must  select  some  one  element  as 
a  standard  and  determine  what  weight  of  every  other  element  will 
combine  with  some  fixed  weight  of  this  standard.  Any  element  might 


THE  LAWS  OF  CHEMICAL  COMBINATION  89 

be  taken  as  the  standard,  and  various  choices  have  been  made  at  dif- 
ferent times.  Convenience  is  the  chief  guide,  and  this  has  finally  fixed 
upon  oxygen  as  the  choice.  Any  value  might  be  assigned  to  oxygen, 
since  the  combining  numbers  are  ratio  numbers  relative  to  the  arbi- 
trary value  assigned  to  the  standard.  The  values  100,  10,  16,  and  8 
have  each  been  taken. as  the  standard  value  of  oxygen,  each  suggested 
by  some  reason  of  convenience.  It  is  natural  that  some  integer  should 
be  chosen,  and  8  is  the  smallest  convenient  one.  Any  smaller  value 
would  make  the  combining  weight  of  hydrogen  less  than  unity,  and 
while  this  would  be  no  serious  matter,  it  is  well  to  have  all  values 
at  least  as  great  as  unity.  We  may  therefore  define  the  combining 
weight  or  the  equivalent  of  an  element  as  that  weight  which  will 
combine  with  8  g.  of  oxygen. 

2.  Experimental  determination  of  the  equivalent.  The  actual  experi- 
mental determination  presents  many  difficulties  when  great  precision 
is  desired.  These  are  due  to  the  various  physical  and  chemical  prop- 
erties of  the  substances  which  must  be  collected  and  weighed,  the 
difficulty  of  obtaining  them  in  an  absolutely  pure  condition,  and  the 
inevitable  loss  in  all  such  operations.  Some  .elements  will  not  combine 
directly  with  oxygen,  and  their  combining  weights  must  be  calculated. 
Thus  with  bromine  we  may  determine  the  ratio  in  which  bromine 
combines  with  silver,  then  the  ratio  of  silver  to  oxygen,  and  from 
these  two  values  calculate  the  ratio  of  bromine  to  oxygen.  Of  course, 
all  the  errors  of  the  two  experimental  ratios  may  accumulate  in  the  cal- 
culated one  and  make  it  less  trustworthy  than  a  direct  determination 
would  be. 

Beginning  with  Berzelius,  who  devoted  many  years  of  his  life  to 
the  determination  of  combining  weights,  some  of  the  most  illustrious 
chemists  of  each  generation  have  exercised  their  greatest  skill  on  this 
problem.  For  many  years  the  work  of  Stas,  which  appeared  largely  be- 
tween the  years  1860  and  1870,  and  which  included  the  determination 
of  the  combining  weights  of  about  a  dozen  of  the  most  common  ele- 
ments, has  been  regarded  as  a  model  of  accuracy.  Some  of  his  equiva- 
lents are  undoubtedly  correct  to  within  a  few  hundredths  of  one  per 
cent.  In  1895  the  American  chemist  Morley  published  the  results  of 
years  of  labor  on  the  combining  ratio  between  oxygen  and  hydrogen, 
the  values  being  trustworthy  to  one  or  two  units  in  the  third  decimal 
place.  At  the  present  time  much  work  is  being  done  along  this  line, 
as  the  combining  weights  of  many  of  the  elements  are  not  accurately 


90  GENERAL  CHEMISTRY 


known,  and  there  are  few  which  cannot  be  more  carefully  determined. 
Theodore  Richards  is  a  leader  in  this  work.  He  has  redetermined  the 
equivalents  of  a  considerable  number  of  the  more  common  elements 
with  a  degree  of  accuracy  never  before  attained  in  any  considerable 
series  of  elements,  and  has  shown  that  even  the  work  of  Stas  requires 
revision  in  view  of  the  greatly  improved  methods  of  research  now 
available. 

3.  Atomic  weights.  For  certain  theoretical  reasons  shortly  to  be  ex- 
plained, some  integer  multiple  of  the  combining  weight  is  more  fre- 
quently employed  than  the  latter  weight  itself.  These  multiples  are 
called  the  atomic  weights,  and  the  table  on  page  12  gives  a  list  of 
them.  The  fact  that  many  are  multiples  of  the  combining  weights  in 
no  way  changes  their  experimental  character  or  modifies  their  funda- 
mental meaning,  for  we  have  seen  that  the  law  of  multiple  proportion 
states  that  a  multiple  of  the  simplest  combining  weight  will  often  be 
necessary  to  express  the  composition  of  a  substance.  The  atomic 
weights  are  merely  those  multiples  which,  for  various  reasons,  afford 
the  greatest  convenience  and  lead  to  the  most  concordant  results. 

III.  The  use  of  atomic  weights  in  expressing  facts.  The  fact  that 
each  element  has  a  distinct  value  of  its  own  in  chemical  reactions,  and 
that  this  value  can  be  experimentally  determined,  suggests  a  number 
of  convenient  applications  of  the  atomic  weights  in  describing  chemical 
changes.  Some  of  these  applications  will  now  be  considered. 

1.  Symbols.    In  the  table  of  atomic  weights  just  referred  to  it  will 
be  noted  that  the  name  of  each  element  is  followed  by  an  abbrevia- 
tion.   This  is  called  its  symbol.    Many  of  these  symbols  are  abbre- 
viations of  the  old  Latin  names  and  bear  no  relation  to  the  current 
English  names.    Thus  sodium  is  designated  as  Na  =  natrium ;  anti- 
mony, as  Sb  =  stibium;  W  stands  for  tungsten,  the  German  name 
for  which  is  wolfram. 

Symbol  weights.  Not  only  does  a  symbol  stand  for  a  certain  ele- 
ment, but  it  also  represents  a  definite  weight  of  it,  namely,  a  weight 
proportional  to  its  atomic  weight.  If,  as  is  usually  the  case,  we  em- 
ploy the  gram  as  our  standard  of  weight,  then  the  symbol  Na  indicates 
23.00  g.  of  sodium;  Sb,  120.2  g.  of  antimony.  Such  a  weight,  ex- 
pressed in  grams,  is  called  the  symbol  weight  of  an  element. 

2.  Formulas.    In  representing  the  composition  of  a  compound  we 
might   disregard   the    symbol   weight   and   use   the    symbols,   thus: 
Fe,  63.52  per  cent ;  S,  36.48  per  cent.    But  obviously  this  would  be 


THE  LAWS  OF  CHEMICAL  COMBINATION  91 

very  awkward.  Instead  of  expressing  composition  in  percentages  it  is 
more  convenient  to  express  it  by  stating  the  number  of  symbol 
weights  of  each  element  present  in  a  combining  weight  of  the  com- 
pound. Thus  the  symbols  FeS  represent  a  compound  made  up  of 
one  symbol  weight  of  iron  and  one  of  sulfur,  namely,  55.84  parts  of 
iron  to  32.07  parts  of  sulfur,  making  a  total  of  87.91  parts  of  iron 
sulfide.  In  any  other  weight  of  iron  sulfide  the  ratio  of  iron  to  sulfur 
will  be  55.84 :  32.07.  To  convert  such  weights  into  percentages, 
remembering  that  the  latter  means  parts  per  hundred,  we  need  merely 
solve  the  proportions : 

87.91 :  55.84  : :  100  :  x  =  63.52  per  cent  iron 
87.91 :  32.07  : :  100  :  x  =  36.48  per  cent  sulfur 

Symbols  used  in  this  way  constitute  the  formula  of  a  compound. 
FeS  is  the  formula  of  iron  sulfide. 

Formula  weights.  Since  each  symbol  in  the  formula  of  a  compound 
represents  a  definite  weight  of  an  element,  it  is  evident  that  the 
formula  must  represent  a  definite  weight  of  the  compound.  This 
weight  is  equal  to  the  sum  of  the  symbol  weights  in  the  formula, 
and  is  called  the  formula  weight  of  the  compound.  Thus  the  formula 
weight  of  iron  sulfide  equals  55.84  +  32.07  =  87.91  g. 

3.  The  calculation  of  formulas  from  percentages.  The  results  of  the 
analysis  of  a  new  compound  are  usually  first  expressed  in  percentages. 
An  example  will  make  clear  the  method  of  calculating  the  formula 
from  percentage  figures.  Suppose  that  an  analysis  of  a  certain  sub- 
stance shows  it  to  contain  31.91  per  cent  K,  28.93  per  cent  Cl,  and 
39.16  per  cent  O.  If  each  of  these  percentages  is  divided  in  turn  by 
the  symbol  weight  of  the  element  it  represents,  we  shall  have  the 
fraction  of  the  symbol  weight  of  each  element  present  in  100  parts  of 

the  compound: 

• 

31.91-^39.10  =  0.8161;    28.92-^-35.46  =  0.8159;    39.16-^16  =  2.447 

These  three  fractional  weights  must  stand  in  the  ratio  of  integer  multi- 
ples of  the  whole  symbol  weights,  so  if  we  divide  all  three  by  the 
smallest,  we  shall  get  the  integer  by  which  the  symbol  weight  must  be 
multiplied  in  each  case.  Dividing  each  by  0.8159,  we  get  the  integers 
1,  1,  3.  The  formula  of  the  compound  is  therefore  KClOg.  Since 
analyses  are  always  slightly  inaccurate,  the  integer  ratios  will  in 
general  differ  slightly  from  true  integers,  but  the  values  will  be  so 


92  GENERAL  CHEMISTRY 

close  as  to  leave  no  doubt  as  to  the  integer  in  question.  The  for- 
mula KC1O3  represents  the  ratio  by  weight  of  the  three  elements  as 
39.1 :  35.46  :  48  in  a  total  of  122.56,  whereas  percentages  represent 
it  as  31.91 :  28.93  :  39.16  in  a  total  of  100.  It  will  be  seen  that  the 
ratios  are  identical. 

We  have  seen  that  water  is  composed  of  11.19  per  cent  hydrogen 
and  88.81  per  cent  oxygen: 

11.19  -*- 1.008  =  11.10  ;  88.81  -s- 16  =  5.55 

The  numbers  11.10  and  5.5  are  in  the  same  ratio  as  are  the  number 
of  symbol  weights  of  hydrogen  and  oxygen  in  water.  This  ratio  is 
2:1;  therefore  the  formula  of  water  is  H2O. 

4.  Equations.    Not  only  may  the  composition  of  compounds  be  repre- 
sented by  formulas,  but  the  changes  taking  place  in  chemical  reactions 
may  be  represented  in  the  form  of  equations.    For  example,  the  elec- 
tric current  decomposes  water  into  hydrogen  and  oxygen.    This  may 
be  represented  by  the  equation 

H2O  =  2  H  +  O 

To  complete  the  proof  that  this  equation  really  represents  the  whole 
truth  we  should  have  to  prove  that  nothing  but  hydrogen  and  oxygen 
is  formed  in  the  reaction,  and  that  18.016  g.  of  water  gives  2.016  g.  of 
hydrogen  and  16  g.  of  oxygen,  as  required  by  the  equation.  It  will 
thus  be  seen  that  an  equation  is  not  merely  algebraic.  The  equation 

HgO  -  Hg  +  O 

represents  the  change  which  takes  place  when  mercuric  oxide  is  heated, 
as  shown  by  experiment.  The  similar  equation 

CaO  =  Ca  +  O 

is  just  as  true  as  an  algebraic  equation,  but  it  is  not  true  chemically, 
since  calcium  oxide  cannot  be  decomposed  in  this  way. 

5.  Representation  of  energy  changes.    We  have  seen  in  Chapter  I  that 
the  changes  in  energy  during  a  chemical  reaction  usually  result  in  the 
liberation  or  absorption  of  heat.    It  is  now  easy  to  adopt  a  system  of 
expressing  these  energy  changes.    We  can  add  to  our  chemical  equa- 
tion the  number  of  heat  units  evolved  or  absorbed  when  the  weight 
of  material  indicated  in  the  equation  undergoes  reaction.  Thus  the 
equation  S  +  2  O  =  SO.  +  71,100  caL 


THE  LAWS  OF  CHEMICAL  COMBINATION  93 

means  that  when  32.07  g.  of  sulphur  is  burned  in  oxygen,  forming 
64.07  g.  of  oxide  of  sulfur,  71,100  cal.  of  heat  are  set  free.  The 
equation  H  +  I  =  HI  -  6036  cal. 

means  that  when  1.01  g.  of  hydrogen  combines  with  126.97  g.  of  iodine 
(in  solid  form),  forming  127.98  g.  of  hydrogen  iodide,  6036  caL  of 
heat  are  absorbed  and  must  be  supplied  from  without  to  maintain  the 
temperature. 

Finally,  it  must  be  noted  that  these  equations  do  not  in  any  way 
describe  the  conditions  under  which  the  reaction  takes  place.  The 
equation  , %. .-  .  Fe  +  S  =  FeS  +  24,000  cal.  '  ' 

affords  no  indication  that  the  reaction  will  not  start  at  ordinary  tem- 
peratures, though  the  large  quantity  of  heat  evolved  suggests  that, 
when  once  started,  it  will  go  on  of  itself.  The  equation 

'  H2O  =  2  H  4-  O  -  69,000  cal. 

does  not  tell  us  how  water  may  be  decomposed,  but  merely  that  in 
decomposition  18.016  g.  will  give  2.016  g.  of  hydrogen  and  16  g.  of 
oxygen,  and  that  a  great  deal  of  energy  (equal  to  69,000  cal.)  is  ab- 
sorbed in  the  reaction.  Experiment  shows  that  this  is  best  supplied 
not  as  heat  but  in  the  form  of  electrical  energy. 

IV.  The  atomic  theory.  The  four  laws  explained  at  length  in  the 
earlier  pages  of  this  chapter  describe  the  chief  characteristics  of  all 
chemical  action  so  far  as  matter  is  concerned.  They  state  such  strik- 
ing peculiarities  that  the  mind  instinctively  seeks  a  theory  which  shall 
coordinate  these  facts,  affording  a  mechanical  picture  of  chemical 
action  in  harmony  with  them.  The  theory  commonly  accepted  is 
known  as  the  atomic  theory,  and  its  essential  features  were  devised 
by  Dalton,  who,  as  we  have  seen,  .first  announced  the  law  of  multiple 
proportion. 

The  essential  features  of  the  atomic  theory.  The  main  features  of 
the  atomic  theory  in  its  present  form,  together  with  the  experimen- 
tal reasons  for  adopting  them,  may  be  very  briefly  stated. 

(1)  It  is  assumed  that  every  weighable  quantity  of  an  elementary 
substance  is  made  up  of  a  very  great  number  of  unit  bodies  which 
Dalton  named  atoms. 

(2)  Since  experiment  shows  that  there  is  no  change  in  mass  when 
two  substances  act  upon  each  other,  it  must  be  true  that  the  masses 


94  GEKEKAL  CHEMISTEY 

of  the   individual    atoms    of  which  these   substances   are   composed 
undergo  no  change  during  the  action. 

(3)  Experiment  also  shows  that  the  composition  of  a  given  com- 
pound is  always  the  same.    The  simplest  way  to  fit  this  fact  into  our 
theory  is  to  assume  that  the  atoms  of  each  element  all  have  the  same 
mass,  while  those  of  different  elements  have  different  masses ;   and 
that  when  one  element  combines  with  another,  the  combination  takes 
place  between  a  definite  number  of  each  kind  of  atoms.    It  will  be 
seen  that  if  these   assumptions  should  be  true,  a  given  compound 
would  of  necessity  have  a  perfectly  definite  composition. 

(4)  The  law  of  multiple  proportions  reminds  us  that  two  elements 
may  combine  in  more  than  one  ratio  to  form  entirely  different  com- 
pounds, each  with  its  own  unchanging  composition.    The  picture  we 
are  forming  can  easily  provide  for  this  peculiarity  if  we  assume  that 
the  two  kinds  of  atoms  can  unite  in  different  ratios.    For  example,  if 
one  atom  of  A  unites  with  one  of  B  under  one  set  of  conditions,  but 
with  two  of  B  under  other  conditions,  then  we  shall  have  the  two 
compounds  AB  and  ABZ.    The  masses  of  B  in  these  two,  combined 
with  a  fixed  mass  of  A,  would  then  be  in  the  ratio  1  :  2,  which  is  in 
complete  accord  with  the  law. 

(5)  The  law  of  combining  weights  tells  us  that  to  each  element 
can  be  assigned  a  number  which  expresses  its  combining  value.    If  we 
assume  that  each  kind  of  atom  has  its  own  peculiar  mass,  and  that  the 
atoms  always  combine  with  each  other  in  definite  numbers,  then  these 
combining  numbers  indicate  the  relative  masses  of  the  atoms  them- 
selves.   The  fact  that  an  element  may  have  two  combining  numbers, 
one  an  integer  multiple  of  the  other,  is  provided  for  by  the  supposi- 
tion that  the  atoms  are  able  to  combine  in  several  different  ratios. 

Summary.  The  picture  of  the  make-up  of  matter  and  the  nature  of 
chemical  action  which  the  atomic  theory  presents  may  be  summed  up 
briefly  as  follows ;  All  matter  is  made  up  of  minute  bodies  called 
atoms.  The  atoms  of  each  element  are  all  alike  in  mass,  but  those  of 
different  elements  have  different  masses.  When  elements  act  upon 
each  other,  the  action  takes  place  between  the  different  kinds  of 
atoms  and  in  definite  numerical  ratios. 

Further  explanations.  To  prevent  misconceptions  some  further  ex- 
planations and  cautions  are  desirable. 

The  size  of  the  atoms.  Dalton  had  no  idea  as  to  the  size  of  these 
atoms,  but  modern  science  has  thrown  much  light  upon  the  question. 


THE  LAWS  OF  CHEMICAL  COMBINATION  95 

Measurements  of  the  thickness  of  soap-bubble  films,  and  of  the  extent 
to  which  a  highly  colored  dye  can  be  diluted  with  water  and  yet  have 
a  perceptible  color,  have  given  us  an  idea  as  to  the  largest  size  which 
it  is  possible  for  these  atoms  to  have.  Lord  Kelvin  has  calculated 
that  it  would  take  4,000,000  hydrogen  atoms,  placed  side  by  side,  to 
make  a  row  1  mm.  long,  and  that  if  a  drop  of  water  were  magnified 
to  the  size  of  the  world,  the  atoms  would  then  be  somewhere  between 
small  shot  and  a  baseball  in  size.  Some  intensely  colored  dyes  give  a 
perceptible  color  to  water  when  1  part  is  dissolved  in  100,000,000 
parts  of  water. 

The  changeability  of  atoms.  It  should  be  noted  that  it  is  not  a  part 
of  the  theory  that  there  are  no  conditions  under  which  atoms  might 
be  changed  in  mass  —  become  transformed  into  other  atoms  or  con- 
verted into  fragments.  It  merely  assumes  that  such  changes  are  not 
a  characteristic  of  chemical  action.  We  shall  see  later  that  there  is 
good  reason  for  thinking  that  some  kinds  of  atoms  do  undergo  trans- 
formations which  change  their  mass  and  character. 

The  different  masses  of  different  kinds  of  atoms.  Experimentally  no 
two  elements  have  been  found  to  have  exactly  the  same  combining 
weight,  though  nickel  (58.68)  and  cobalt  (58.97)  approach  each 
other  very  closely.  Therefore  we  assume  that  each  kind  of  atom  has 
its  own  peculiar  mass.  It  is  quite  conceivable,  however,  that  the  atoms 
of  two  elements  should  have  exactly  the  same  mass,  the  difference 
between  them  lying  in  other  properties,  such  as  energy  content. 

The  mass  of  each  atom  of  the  same  element  not  necessarily  the  same. 
A  little  reflection  will  show  that  the  atoms  of  the  same  element  may 
not  have  precisely  the  same  mass.  In  every  wei'ghable  quantity  of  an 
elementary  body  there  must  be  millions  of  atoms,  so  that  all  we  can 
infer  is  that  the  average  mass  is  always  the  same,  within  our  ability  to 
determine  it.  Now  the  mass  of  a  man  as  determined  from  the  average 
of  a  group  of  three  men  selected  at  random  will  vary  greatly;  that 
from  a  group  of  three  hundred  men  will  vary  less ;  that  from  a  group  of 
three  thousand,  still  less.  As  we  increase  the  number  from  which  the 
average  is  calculated,  we  come  to  a  place  where  variations  can  no 
longer  be  detected.  On  the  other  hand,  the  atoms  of  an  element 
cannot  differ  much  from  the  average,  or  we  should  be  able  to  separate 
the  lighter  individuals  from  the  heavier  in  various  ways.  The  state- 
ment that  the  atoms  of  an  element  all  have  the  same  weight,  there- 
fore, means  that  the  differences  in  weight,  if  any,  cannot  be  detected. 


96  GEKEBAL  CHEMISTEY 

Molecules  and  atoms.  In  developing  his  atomic  theory  Dalton  made 
no  distinction  between  elements  and  compounds,  each  being  regarded 
as  composed  of  atoms.  Evidently  the  smallest  particle  of  a  com- 
pound must  be  made  up  of  several  different  kinds  of  atoms,  while  the 
smallest  particle  of  a  gaseous  element,  like  oxygen,  might  consist 
either  of  individual  atoms  or  of  small  groups  of  similar  atoms  combined 
like  the  atoms  in  a  compound  particle.  In  time  it  was  found  conven- 
ient to  make  a  distinction.  The  term  molecule  is  now  applied  to  the 
smallest  unit  either  of  a  compound  or  of  an  element,  which,  taken 
in  large  aggregations,  makes  up  the  bodies  we  deal  with,  and  into 
which  substances  can  be  divided  without  chemical  decomposition. 
The  term  atom  is  applied  to  the  smallest  unit  of  an  element  which 
takes  part  in  a  chemical  reaction. 

Atomic  weights  and  molecular  weights.  Since  the  mass  of  each  kind 
of  atom  is  always  the  same,  it  will  be  seen  that  some  one  set  of  num- 
bers may  be  chosen  which  will  represent  the  relative  masses  of  each 
of  the  atoms,  and  that  these  numbers  will  be  either  the  combining 
numbers  themselves  or  some  integer  multiple  of  them.  The  choice 
of  the  real  atomic  ratio  number  from  among  the  possible  combining 
numbers  has  presented  great  difficulties  historically,  but  we  now  have 
sound  principles  for  our  guidance.  These  will  be  developed  in  a  sub- 
sequent chapter.  For  the  present  we  may  assume  that  the  numbers 
in  the  atomic  table  (p.  12)  do  represent  the  real  atomic  ratios,  and 
are  therefore  properly  called  atomic  weights.  When  we  make  use  of 
the  atomic  theory,  therefore,  in  expressing  chemical  facts,  the  symbol  of 
an  element  represents  an  atom,  while  a  formula  represents  a  molecule. 

Thus  we  have  good  reason  for  thinking  that  the  substance  we 
call  oxygen  is  made  up  of  oxygen  molecules  having  the  composition 
represented  by  the  formula  O2.  When  water  is  decomposed,  the  first 
action  seems  to  be  the  parting  of  the  water  molecules  into  the  two 
kinds  of  atoms:  H2O  =  2H  +  O 

These  at  once  combine  to  form  molecules  of  the  two  gases,  so  that 
the  final  condition  is  represented  by  the  equation 


With  mercury  and  other  metals,  when  in  the  form  of  vapor,  the 
atoms  do  not  appear  to  form  molecules,  so  that  mercury  vapor  is 
represented  by  the  formula  Hg,  not  Hg  . 


THE  LAWS  OF  CHEMICAL  COMBINATION  97 

Gram-atomic  and  gram-molecular  weights.  In  accordance  with  the 
usage  just  explained  the  term  gram-atomic  weight  is  often  employed 
instead  of  symbol  weight,  and  gram-molecular  weight,  or  molar  weight, 
instead  of  formula  weight.  Both  really  mean  the  same  thing,  namely, 
as  many  grams  of  the  element  or  compound  as  there  are  units  in  its 
symbol  or  formula. 

Molecular  formulas.  If,  as  is  the  custom,  we  make  use  of  formulas 
to  represent  not  only  the  composition  of  a  compound,  but  also  the 
number  of  atoms  in  its  molecule,  it  may  easily  happen  that  the  sim- 
plest formula,  calculated  according  to  the  method  already  described 
(p.  91),  will  not  fulfill  these  conditions.  Thus  the  simplest  formula 
correctly  representing  the  composition  of  hydrogen  peroxide  is  HO ; 
but  there  is  good  reason  for  believing  that  the  molecules  of  the  com- 
pound really  consist  of  two  atoms  of  hydrogen  and  two  of  oxygen, 
giving  the  formula  H2O2.  It  has  been  found  possible  to  devise 
methods  which  give  us  the  true  molecular  formulas  (or  molecular 
weights)  of  most  compounds,  and  these  will  be  described  in  a  sub- 
sequent chapter. 

The  value  of  the  atomic  theory.  Like  any  good  theory,  the  atomic 
theory  is  useful  chiefly  in  affording  a  concrete  image  of  the  mechanics 
of  the  topic  it  covers  —  in  this  case  chemical  action  —  and  in  sug- 
gesting profitable  investigation.  It  does  not  make  a  great  deal  of 
difference  whether  or  not  such  things  as  atoms  actually  exist,  and 
nothing  short  of  seeing  them  will  ever  conclusively  settle  the  question. 
The  main  question  is,  Does  the  mechanical  conception  correspond  to 
all  known  facts  ?  If  not,  what  are  the  exceptions,  so  that  we  may 
keep  them  in  mind  and  not  be  misled  by  the  theory?  It  may  be 
said  that  the  theory  has  been  found  to  be  in  remarkable  accord  with 
all  we  know  about  matter,  that  it  has  been  the  incentive  to  a  vast 
amount  of  profitable  work,  and  that  the  number  of  facts  not  in  accord 
with  the  theory  is  not  large.  When  properly  regarded  it  is  a  very 
helpful  conception. 


CHAPTER  VIII 

EQUATIONS  AND  CALCULATIONS 

In -the  chapters  on  Oxygen,  Hydrogen,  and  Water,  reactions  were 
described  which  must  now  be  reviewed  and  put  into  the  form  of 
definite  equations.  The  student  will  find  it  profitable  to  turn  back 
to  the  description  and  associate  the  reactions  with  the  equations  as 
they  are  now  given. 

1.  The  preparation  of  oxygen. 

Decomposition  of  mercuric  oxide  :   2  HgO  =  2  Hg  -f  O2  (1) 

Decomposition  of  potassium  chlorate :   2  KC1O3  =  2  KC1  +  3O2  (2) 

Decomposition  of  manganese  dioxide :  3  MnO2  =  MngO4  +  O2  (3) 

Decomposition  of  barium  peroxide  :  2  BaO2  =  2  BaO  +  O2  (4) 
Action  of  sodium  peroxide  on  water : 

2  Na202  +  2  H20  =  4  NaOH  +  O2  (5) 

Electrolysis  of  water :  2  H2O  =  2  H2  +  O2  (6) 

It  must  be  constantly  kept  in  mind  that  such  equations  are  merely 
a  very  condensed  way  of  stating  quantitative  chemical  facts.  In  read- 
ing the  equation  the  full  statement  should  be  recalled.  Thus  equa- 
tion (2)  is  a  statement  that  under  appropriate  conditions  two  formula 
weights  of  potassium  chlorate  give  two  formula  weights  of  potassium 
chloride  and  three  of  oxygen.  In  terms  of  the  atomic  theory  the  state- 
ment is  that  two  molecules  of  the  chlorate  give  two  of  the  chloride 
and  three  of  oxygen. 

2.  The  action  of  oxygen  upon  elements. 

Combustion  of  sulfur:  S  -f  O2  =  SO2  (7) 

Combustion  of  phosphorus :  4  P  +  5  O2  =  2  P2O5  (8) 

Combustion  of  sodium :   2  Na  +  O2  =  Na2O2  (9) 

Combustion  of  iron :  3  Fe  +  2  O2  =  FegO4  (10) 

It  will  be  noted  that  these  four  elements  are  represented  as  though 
their  molecules  consisted  of  but  a  single  atom.  This  is  probably  not  the 
case,  but  we  have  no  satisfactory  way  of  determining  how  many  atoms 
are  present  in  the  molecule  of  any  solid  substance,  and  in  the  absence 

98 


EQUATIONS  AND  CALCULATIONS  99 

of  other  information  we  represent  it  by  the  simplest  possible  formula. 
On  the  other  hand,  the  solid  compound  sodium  peroxide  (Na2O2) 
might  be  represented  by  the  formula  NaO  just  as  well,  but  the  double 
formula  more  satisfactorily  expresses  many  of  its  reactions. 

3.  The  preparation  of  hydrogen. 

Electrolysis  of  water :  2  H2O  =  2  H2  +  O2  (11) 

Decomposition  of  water  by  sodium : 

2  H2O  +  2  Na  =  2  NaOH  +  H2  (12) 
Decomposition  of  water  by  magnesium : 

2  H20  +  Mg  =  Mg(OH)2  +  H2  (1 3) 
Decomposition  of  water  by  iron : 

3  Fe  +  4  H20  =  Fe3O4  +  4  H2  (14) 
Action  of  acids  on  metals. 

Iron  4-  hydrochloric  acid :  Fe  4-  2  HC1  =  FeCl2  +  H2  (15) 

Zinc  4-  hydrochloric  acid :  Zii  +  2  HC1  =  ZnCl2  +  H2  (16) 

Iron  +  sulfuric  acid :  Fe  +  H2SO4  =  FeSO4  +  H2  (17) 

Zinc  4-  sulfuric  acid :  Zn  +  H2SO4  =  ZnSO4  +  H2  (18) 

Zinc  +  sodium  hydroxide  :  Zn  +  2  NaOH  =  Zn(ONa)2  +  H2  (19) 

The  formula  Mg(OH)2  in  equation  (13)  is  merely  a  more  convenient 
way  of  writing  MgO2H2,  and  represents  the  same  thing. 

4.  The  action  of  hydrogen  with  other  substances. 

Hydrogen  +  chlorine  :  H2  +  C12=2HC1  (20) 

Hydrogen  +  sulfur :  H2  +  S  =  H2S  (21) 

Hydrogen  +  nitrogen  :  3H2  +  N2=2NH3  (22) 

Hydrogen  +  oxygen :  2H2  +  O2=2H2O  (23) 

Hydrogen  +  copper  oxide :  H2  +  CuO  =  Cu  +  H2O  (24) 

Hydrogen  4-  iron  oxide  :  4  H2  4-  Fe3O4  =  3  Fe  4-  4  HaO  (25) 

5.  Water  of  crystallization.    Experiment  has  shown  that  the  water 
with  which  certain  compounds  unite  in  forming  crystals  always  bears 
a  definite  ratio  by  weight  to  the  compound,  and  the  resulting  hydrate 
should  therefore  have  a  definite  formula.    Thus  the  blue  hydrate  of 
copper  sulfate,  when  heated,  loses  36.07  per  cent  of  its  weight  as 
water,  which  corresponds  to  five  formula  weights  of  water  to  one  of 
the  sulfate.    The  anhydrous  copper  sulfate  has  the  formula  CuSO4, 
so  that  the  composition  of  the  hydrate  may  be  expressed  by  the  for- 
mula CuSO9H1Q.    It  is  customary,  however,  to  retain  the  formulas  of 


100  GENERAL   CHEMISTRY 

the  two  original  compounds,  indicating  the  hydrate  by  the  formula 
CuSO4-5H2O.  When  this  hydrate  is  heated,  the  hydrogen  and 
oxygen  are  given  off  as  water,  leaving  the  copper  sulf ate : 

CuSO4  •  5  H2O  =  CuSO4  +  5  H2O  (26) 

It  must  be  understood  that  this  usage  is  for  convenience  in  recogniz- 
ing that  the  compound  is  a  hydrate  of  copper  sulfate,  and  is  not 
intended  to  convey  the  impression  that  water  as  such  exists  in  the 
hydrate. 

6.   The  preparation  of  hydrogen  peroxide. 

Barium  peroxide  +  sulfuric  acid : 

Ba02  +  H2S04  -  BaS04  +  H2O2  (27) 

Barium  peroxide  -f  phosphoric  acid : 

3  Ba02  +  2  H3P04  =  Ba3(PO4)2  +  3  H2O2  (28) 

Types  of  reactions.  A  study  of  the  reactions  so  far  considered  shows 
that  they  may  be  classified  into  four  distinct  types,  and  we  shall  find 
that  almost  all  of  the  reactions  which  we  shall  meet  can  be  assigned 
to  one  of  these  four. 

1.  Direct  union.    Two  elements  (or  compounds)  may  unite  to  form 
a  compound.    This  is  illustrated  by  the  combustion  of  substances  in 
oxygen  and  the  action  of  hydrogen  upon  different  elements.   (Equa- 
tions (7)  to  (10)  and  (20)  to  (23).) 

2.  Decomposition.    A   compound  may   decompose    into   other  com- 
pounds or  elements,  as  is  the  case  when  mercuric  oxide,  potassium 
chlorate,  or  manganese  dioxide  is  heated.    (Equations  (1)  to  (4).) 

3.  Substitution.    One  element  may  take  the  place  of  another  in  a 
compound,  the  substituted  element  being  set  free.    This  is  the  case 
when  the  metals  zinc  and  iron  act  upon  hydrochloric  or  sulfuric  acids, 
liberating  hydrogen.    (Equations  (15)  to  (18).) 

4.  Double  decomposition.    This  is  probably  the  most  common  type  of 
reaction  and  consists  in  the  interchange  of  two  elements  present  in 
two  different  compounds,  thus   resulting  in   the  formation  of  two 
new  compounds.    It  is  illustrated  in  equations  (27)  and  (28),  which 
represent  the  reaction  taking  j^ace  in  the  preparation  of  hydrogen 
peroxide.    In  these  reactions  it  will  be  seen  that  the  barium  of  barium 
peroxide  and  the  hydrogen  of  the  acids  change  places,  forming  two 
new  compounds. 


EQUATIONS  AND  CALCULATIONS  101 

Chemical  calculations.  Equations  such  as  the  ones  which  have  just 
been  presented  are  merely  a  convenient  form  in  which  to  record  the 
results  of  experiment.  They  represent  the  composition  of  substances 
entering  into  reaction  with  each  other,  and  the  proportion  by  weight 
in  which  they  take  part  in  the  reaction ;  they  also  represent  the 
weights  and  compositions  of  the  products  formed.  When  properly 
verified  by  experiment  they  may  be  employed  very  conveniently  in 
solving  many  problems  which  involve  chemical  transformations.  A 
few  typical  examples  will  now  be  given,  together  with  the  method  of 
solution,  and  the  student  should  supplement  these  from  any  one  of 
the  problem  books  now  available.1 

1.  What  weight  of  oxygen  can  be  obtained  by  heating  20  g.  of 
potassium  chlorate  ? 

The  formula  weight  (or  molecular  weight)  of  potassium  chlorate  is 

39.1  +  35.46  +  3  x  16  =  122.56 
Equation  (2) :          2  KC1O3  =  2  KC1  +  3  O5 


2 
2  x  122.56  3  X  32 


The  equation  tells  us  that  the  ratio  by  weight  between  the  chlorate 
decomposed  and  the  oxygen  obtained  is  always  245.12 :  96.  The 
quantity  obtainable  from  20  g.  of  chlorate  may  therefore  be  found 
from  the  proportion: 

245.12  :  96  :  :  20  :  x 
x  =  7.83  g. 

2.  What  weight  of  oxygen  will  be  required  for  the  combustion  of 
25  g.  of  iron  ? 

Equation  (10) :  3  Fe  +  2  O2  =  Fe3O4 

3  X  55.84         2  X  32 

The  iron  stands  to  the  oxygen  in  the  ratio  167.52  :  64.  Consequently 

167.52  :  64  :  :  25  :  x 
z=9.55g. 

3.  What  weight  of  zinc  will  be  required  to  liberate  10  g.  of  hydrogen 
by  its  action  with  hydrochloric  acid  ? 

Equation  (16) :        Zn  +  2  HC1  =  ZnCl2  +  H2 

65.37  72.94  136.S  2.02 

Since  hydrogen  and  zinc  stand  in  the  ratio  2.02 :  65.37,  therefore 

2^02  :  65.37  :  :  10  ix 
x  =  323.6  g.  zinc 

1  Hale,  Calculations  of  General  Chemistry. 


102  GENERAL  CHEMISTRY 

4.  What  weight  of  zinc  chloride  (ZnCl2)  will  be  produced  at  the 
same  time? 

The  hydrogen  and  zinc  chloride  are  formed  in  the  ratio  2.02  :  136.3  ; 
consequently  2.02  :  136.3  ::  10  :  z 

z=  674.3 

5.  Some  of  the  common  acids  are  always  sold  in  the  form  of  con- 
centrated solutions,  the  label  on  the  bottle  bearing  a  statement  as  to 
the  density  of  the  solution  and  the  percentage  by  weight  of  acid  in  it. 
Thus  the  concentrated  hydrochloric  acid  of  commerce  usually  has  a 
density  of  1.20  and  contains  40  per  cent  of  the  compound  HC1.    Let 
it  be  required  to  calculate  what  volume  of  such  a  solution  will  be  used 
in  producing  10  g.  of  hydrogen,  as  in  problem  3.    It  must  be  remem- 
bered that  the  volume  of  a  liquid  multiplied  by  its  density  gives  its 
weight  (that  is,  Fx  D  =  W£),  and  that  of  this  weight  only  40  per  cent 
is  the  material  we  are  concerned  with.    The  weight  of  pure  hydrogen 
chloride  (HC1)  is  obtained  from  the  proportion 

2.02  :  72.94  :  :  10  :  x 
z=361g.  HC1 


This  is  40  per  cent  of  the  weight  of  the  solution  required,  the  whole 

361 
solution  weighing  —  -  x  100  =  902.5  g.    Dividing  this  by  the  density 

of  the  solution,  1.2  (in  other  words,  by  the  weight  of  1  cc.),  we  have 


Problems  involving  calculations  of  volume.  It  is  often  required  to 
calculate  the  volume  of  a  gas  which  can  be  obtained  in  a  given  reaction 
when  measured  under  ordinary  laboratory  conditions,  or  what  weight 
of  materials  will  be  required  to  produce  a  desired  volume.  Since 
chemical  equations  deal  with  weights,  not  volumes,  it  is  always  neces- 
sary, as  a  first  step,  to  determine  the  weight  of  the  gas,  and  from  this 
to  calculate  its  volume.  Since  the  weight  of  a  unit  volume  of  a  gas 
varies  with  every  change  in  temperature  and  pressure,  it  would  evi- 
dently be  a  great  task  to  tabulate  all  possible  weights  for  reference. 
Consequently  the  tables  (such  as  the  one  in  the  Appendix)  give 
merely  the  weight  under  standard  conditions,  and  a  measured  vol- 
ume of  a  gas  must  always  be  reduced  to  these  conditions  before  its 
weight  can  be  calculated. 


EQUATIONS  AND  CALCULATIONS  103 

6.  What  volume  of  oxygen,  measured  under  ordinary  laboratory 
conditions  (say  750  mm.  and  20°),  may  be  obtained  by  heating  100  g. 
of  mercuric  oxide  ? 

Equation  (1) :  2  HgO  =  2  Hg  +  O2 

2  X  216.6  32 

433.2  :  32  : :  100  :  x 
#  =  7.4  g.  of  oxygen 

1  1.  of  oxygen  (standard)  weighs  1.429  g.  (see  Appendix) 
7.4  -f- 1.429  =  5.18  1.  (standard) 

To  ascertain  the  volume  which  this  will  occupy  under  laboratory 
conditions  we  may  employ  the  equation  on  page  49 : 

Px  Fx273 

760  XT 
Substituting,  we  have 

750  x  Fx273 


5.18  = 


760  x  293 


Solving  for  F,  the  volume  under  laboratory  conditions,  we  obtain 
Y=  5.631. 

7.  How  many  grams  of  potassium  chlorate  will  be  required  to 
yield  101.  of  oxygen  measured  under  ordinary  laboratory  conditions  ? 
In  this  case  the  volume  under  standard  conditions  must  be  calculated 
before  the  weight  can  be  found. 

PxFx273      750x10x273      Q1Q1 
F<  =      760  XT     =      760  x  293      =  9'19  L  (standard) 

9.19  x  1.429  =  13.13  gl  oxygen 
Equation  (2)  •  2  KC1O3  =  2  KC1  +  3O9 

245.12  96 

96  :  245.12  :  :  13.13  :  x 
x  =  33.53  g.  potassium  chlorate 

It  is  useful  to  observe,  as  a  check  on  such  calculations,  that  in 
reducing  to  standard  conditions  a  volume  measured  under  ordinary 
laboratory  conditions,  the  decrease  in  volume  amounts  to  about  8  or 
10  per  cent,  with  a  corresponding  increase  when  the  reduction  is  in 
the  reverse  direction.  The  student  should  also  cultivate  the  habit  of 
noticing  whether  his  results  seem  reasonable,  since  arithmetical  errors 
are  easily  made  in  such  calculations. 


CHAPTER  IX 

NITROGEN  AND  THE  RARE  ELEMENTS:   HELIUM,  NEON,  ARGON, 

KRYPTON,  XENON 

Historical.  The  attention  of  investigators  was  first  drawn  to  the 
general  subject  of  gases  by  Joseph  Black  (1728-1799),  professor  of 
chemistry  in  the  University  of  Edinburgh,  who  discovered  the  gas  now 
known  as  carbon  dioxide  and  studied  its  properties.  Following  the 
work  of  Black  we  have  the  investigations  of  Priestley,  Cavendish, 
Scheele,  Lavoisier,  and  others,  which  resulted  in  the  discovery  .of 
oxygen  and  hydrogen  and  in  the  determination  of  the  composition  of 
water.  It  was  during  this  period  (1772)  that  Rutherford,  professor 
of  botany  in  the  University  of  Edinburgh,  prepared  nitrogen  and  rec- 
ognized it  as  a  new  substance,  regarding  it,  however,  as  a  combination 
of  air  with  phlogiston.  A  little  later  Scheele  pointed  out  that  this 
new  substance  is  a  normal  constituent  of  the  air.  Lavoisier  was  the 
first  to  regard  it  as  an  element,  and  gave  to  it  the  name  azote,  a  word 
signifying  that  it  does  not  support  life.  Later  the  element  was  found 
to  be  an  essential  constituent  of  niter  (potassium  nitrate),  and  because 
of  this  fact  Chaptal  termed  it  nitrogen,  meaning  "  niter-producing." 

Occurrence.  In  the  free  condition  nitrogen  occurs  in  large  quantities 
in  the  atmosphere,  mixed  with  oxygen  and  small  amounts  of  other 
gases.  In  100  volumes  of  dry  air  there  are  approximately  78  vol- 
umes of  nitrogen,  21  of  oxygen,  and  1  of  other  gases.  In  the  combined 
state  nitrogen  is  present  in  sodium  nitrate  (NaNO8)  and  potassium 
nitrate  (KNO3),  both  of  which  occur  in  nature  in  considerable  quan- 
tities. It  is  likewise  an  essential  constituent  of  the  compounds  known 
as  proteins,  which  are  present  in  all  living  organisms.  The  human 
body  contains  about  3  per  cent  of  nitrogen.  The  value  of  organic 
matter  of  various  kinds  as  a  constituent  of  fertilizers  is  based  largely 
on  its  nitrogen  content. 

Preparation  of  nitrogen.  Nitrogen  may  be  obtained  either  from  air 
or  from  compounds  of  nitrogen. 

1.  Preparation  from  air.  Nitrogen  is  ordinarily  obtained  from  air. 
To  separate  it  from  the  oxygen  with  which  it  is  mixed,  the  air  is 

104 


NITROGEN  AND   THE  RARE  ELEMENTS 


105 


brought  in  contact  with  some  substance  which  combines  with  the 
oxygen  but  which  has  no  effect  upon  the  nitrogen.    The  substances 
ordinarily  used  for  withdrawing  the  oxygen  are  phosphorus  and  cop- 
per, not  only  because  they  combine  read- 
ily with   oxygen  but  also  because  the 
oxides  formed  are  solids  and  thus  admit 
of  an  easy  separation  from  the  remaining 
nitrogen. 

The  preparation  of  nitrogen  from  air  through 
the  action  of  phosphorus  is  commonly  conducted 
FIG.  50  as  follows  :  The  phosphorus  is  placed  in  a  small 

porcelain  dish  supported  on  a  cork  and  floated 

on  water  (Fig.  50).  It  is  then  ignited  by  contact  with  a  hot  wire,  and  a  bell  jar 
is  immediately  brought  over  it  so  as  to  confine  a  portion  of  the  air.  The  phos- 
phorus combines  with  the  oxygen  to  form  an  oxide  known  as  phosphorus  pentox- 
ide.  This  is  a  white  solid  which  floats  about  in  the  bell  jar ;  in  a  short  time  it  is 
all  absorbed  by  the  water,  leaving  the  nitrogen.  The  withdrawal  of  the  oxygen 
is  indicated  by  the  rising  of  the  water  in  the  bell  jar. 

When  copper  is  used  for  separating  the  oxygen  from  the  nitrogen,  the  opera- 
tion may  be  conducted  as  follows :  The  metal  is  placed  in  a  tube  A  (Fig.  51) 
and  heated  to  a  high  temperature.  The  air  is  then  forced  slowly  through  the 
tube  by  allowing  a  small  stream  of  water  to  flow  into  a  bottle  B  connected 
with  the  tube,  as  shown  in  the  figure.  The  oxygen  combines  with  the  hot  copper 
and  forms  copper 
oxide  (CuO),  a  solid 
which  remains  in  the 
tube  while  the  nitro- 
gen passes  on  and  is 
collected  over  water 
in  a  cylinder  C. 

Inasmuch  as 
air  contains  small 
percentages  of 
other  gases  be- 
sides oxygen  and 

nitrogen,  and  since  the  phosphorus,  as  well  as  the  copper,  removes 
only  the  oxygen,  it  is  evident  that  the  nitrogen  obtained  by  these 
methods  is  never  quite  pure ;  about  1  per  cent  of  the  product  is 
composed  of  other  gases,  from  which  it  is  very  difficult  to  separate 
the  nitrogen.  The  impure  nitrogen  so  obtained,  however,  may  be  used 
for  a  study  of  most  of  the  properties  of  the  element,  since  these  are 
not  materially  affected  by  the  presence  of  the  other  gases.  The  most 


FIG.  51 


106  GENERAL  CHEMISTRY 

economical  way  of  preparing  pure  nitrogen  consists  in  liberating  it 
from  some  of  its  compounds. 

2.  Preparation  from  compounds.  The  compound  most  commonly  used 
for  the  preparation  of  nitrogen  is  ammonium  nitrite  (NH4NO2).  When 
heated,  this  compound  decomposes  into  nitrogen  and  water,  as  repre- 
sented iji  the  following  equation: 

NH4N02  =  N2  +  2H20 

Since  ammonium  nitrite  is  not  readily  kept  in  the  pure  state,  it  is  con- 
venient to  substitute  for  it  a  mixture  of  sodium  nitrite  (NaNO2)  and 
ammonium  chloride  (NH4C1).  These  two  compounds  react  to  form 
sodium  chloride  and  ammonium  nitrite : 

"NaNO2  +  NH4C1  =  NH4NO2  +  NaCl 

As  fast  as  formed,  the  ammonium  nitrite  decomposes  into  nitrogen 
and  water,  as  explained  above. 

Nitrogen  may  also  be  prepared  from  ammonium  chloride.  When 
this  compound  is  mixed  with  potassium  dichromate  and  the  mixture  is 
heated,  nitrogen  is  evolved  in  accordance  with  the  following  equation : 

K2Cr207  +  2  NH4C1  =  2  KC1  +  Cr2O3  +  4  H2O  +  N2 

Commercial  methods  of  preparation.  It  is  evident  that  the  method 
used  for  the  preparation  of  oxygen  from  liquid  air  would  serve  equally 
well  for  the  preparation  of  nitrogen.  In  addition  to  this  method  it  is 
also  obtained  commercially  by  the  action  of  copper  upon  air,  as  ex- 
plained above.  The  copper  oxide  formed  is  reduced  to  metallic  copper 
(see  pp.  42-43),  which  is  again  used.  Natural  gas  is  frequently  used 
as  the  reducing  agent  in  place  of  hydrogen. 

Properties.  Nitrogen,  like  hydrogen  and  oxygen,  is  a  colorless,  odor- 
less, tasteless  gas.  It  is  slightly  lighter  than  oxygen,  1  1.  of  it  weigh- 
ing 1.2507  g.  Its  solubility  in  water  is  about  the  same  as  that  of 
hydrogen,  1 1.  of  water  dissolving  about  20  cc.  of  the  gas  under 
standard  conditions.  At  its  critical  temperature,  —  146°,  it  is  con- 
densed to  a  liquid  by  a  pressure  of  33  atmospheres.  Liquid  nitrogen 
is  colorless,  boils  at  —  195.7°,  and  has  a  density  of  0.8  at  its  boiling 
point.  At  a  still  lower  temperature  it  may  be  obtained  in  the  form 
of  an  icelike  solid  melting  at  —  210.5°. 

Chemical  conduct.  Nitrogen  is  much  less  active  than  oxygen,  show- 
ing little  or  no  tendency  to  combine  with  any  other  elements  at  ordi- 
nary temperatures.  Nevertheless,  at  high  temperatures  and  under 


NITROGEN  AND^  THE  BARE  ELEMENTS  107 

suitable  conditions  it  combines  with  many  of  the  elements.  Thus, 
when  subjected  to  the  influence  of  electric  sparks,  nitrogen  combines 
with  hydrogen  to  form  ammonia  (NH3),  and  with  oxygen  to  form 
nitric  oxide  (NO).  Nitrogen  likewise  combines  directly  with  silicon, 
boron,  titanium,  and  also  with  most  of  the  metals,  notably  lithium, 
magnesium,  and  calcium.  The  compounds  formed  by  the  union  of 
nitrogen  with  another  element  are  in  general  termed  nitrides,  just  as 
the  compounds  formed  by  the  union  of  oxygen  with  another  element 
are  termed  oxides. 

The  assimilation  of  nitrogen  by  plants.  While  nitrogen  is  an  essen- 
tial constituent  of  both  plants  and  animals,  yet,  with  the  exception  of 
a  few  plants,  especially  those  belonging  to  the  natural  order  Legu- 
minosse,  these  organisms  have  not  the  power  of  directly  assimilating 
free  nitrogen  from  the  atmosphere,  but 
obtain  then-  supply  from  certain  compounds 
of  nitrogen.  It  has  long  been  known  that 
some  of  the  leguminous  plants,  such  as  the 
beans,  peas,  and  clover,  not  only  thrive  in 
poor  soil  but  at  the  same  time  enrich  it. 
Investigation  has  shown  that  these  plants 
obtain  at  least  a  portion  of  their  supply 
of  nitrogen  from  the  atmosphere.  The 

assimilation  of. nitrogen  is  accomplished  through  the  agency  of  groups 
of  microorganisms  which  are  gathered  on  little  tubercles  on  the  roots 
of  the  plants,  as  represented  in  Fig.  52,  which  shows  the  tubercles  on 
the  roots  of  a  variety  of  bean.  These  microorganisms  have  the  power 
of  converting  free  nitrogen  taken  from  the  air  into  compounds  of 
nitrogen,  some  of  which  are  assimilated  by  the  plant,  while  others  are 
left  in  the  soil  and  thus  enrich  it. 

Uses  of  nitrogen.  Free  nitrogen  is  used  to  a  limited  extent  in  the 
preparation  of  certain  nitrogenous  compounds  (p.  424)  employed  as 
fertilizers.  Mercurial  thermometers  designed  for  use  at  temperatures 
of  from  300°  to  500°  are  filled  (over  the  mercury)  with  nitrogen  under 
pressure.  In  this  way  the  mercury  is  prevented  from  boiling,  even  at 
temperatures  considerably  above  its  ordinary  boiling  point  (356.6°). 
The  expansion  of  nitrogen  itself,  according  to  the  law  of  Gay-Lussac, 
is  also  utilized  as  a  measure  of  temperature,  at  points  above  the  range 
of  a  mercurial  thermometer. 


108  GENERAL  CHEMISTKY 

THE    RARE    ELEMENTS    IN   THE    ATMOSPHERE:    HELIUM,  NEON, 
ARGON,  KRYPTON,  XENON 

The  elements  named  above  are  given  in  the  order  of  their  atomic 
weights.  Argon  will  be  considered  first,  since  it  was  the  first  one 
discovered  and  is  the  most  abundant. 

Argon.  Attention  has  been  called  to  the  fact  that  oxygen  and  ni- 
trogen combine  under  the  influence  of  electric  sparks  to  form  nitric 
oxide  (p.  107).  In  1785  Cavendish,  in  the  course  of  his  experi- 
ments on  air,  passed  electric  sparks  through  an  inclosed  volume  of  air 
(nitrogen  and  oxygen),  introducing  more  oxygen  from  time  to  time, 
so  as  to  make  sure  that  the  quantity  of  oxygen  present  would  be 
sufficient  to  combine  with  all  the  nitrogen.  After  repeated  sparking, 
the  oxide  of  nitrogen  formed  by  the  union  of  the  gases,  together  with 
the  excess  of  oxygen,  were  removed  by  absorbing  them  in  appropri- 
ate liquids.  In  this  experiment  Cavendish  observed  that  even  after 
repeated  trials  there  still  remained  a  small  residue  of  gas,  in  volume 
about  y^  of  the  air  taken,  which  would  not  combine  with  oxygen, 
and  therefore  presumably  was  not  nitrogen.  No  attention  was  paid 
to  this  observation  until  1894,  when  Lord  Rayleigh  observed  that  the 
density  of  nitrogen  obtained  from  air  was  about  0.5  per  cent  greater 
than  the  density  of  nitrogen  obtained  from  its  compounds.  After 
repeating  his  experiments  a  number  of  times,  always  with  the  same 
results,  Rayleigh  concluded  that  the  most  reasonable  explanation  of 
the  difference  in  the  densities  of  the  gas  obtained  from  the  two  differ- 
ent sources  lay  in  the  supposition  that  the  nitrogen  which  he  obtained 
from  air  contained  a  small  amount  of  some  gas  heavier  than  nitrogen. 
This  conclusion  recalled  the  experiments  of  Cavendish,  and  it  was 
thought  that  perhaps  the  residue  of  gas  which  Cavendish  obtained  in 
his  experiment  was  the  same  which  caused  the  higher  density  of  the 
atmospheric  nitrogen.  Acting  on  this  assumption,  Rayleigh  and  the 
English  chemist  Ramsay  attempted  to  isolate  any  unknown  gas  which 
might  be  mixed  with  the  atmospheric  nitrogen.  Rayleigh  employed 
the  method  of  Cavendish,  while  Ramsay  attempted  to  separate  the 
unknown  gas  by  repeatedly  passing  the  atmospheric  nitrogen  over 
heated  magnesium,  which  combines  readily  with  nitrogen.  Without 
going  into  detail,  it  is  sufficient  to  state  that  both  investigators  suc- 
ceeded in  showing  that  the  atmospheric  nitrogen,  as  ordinarily  prepared, 
in  reality  contains  a  small  percentage  of  gas  differing  from  nitrogen  in 


NITROGEN  AND  THE  BARE  ELEMENTS  109 

that  it  has  a  greater  density  and  does  not  unite  with  any  other  elements. 
This  gas  proved  to  be  a  new  element,  to  which  the  name  argon  was 
given.  In  its  preparation  in  larger  quantities  it  has  been  found  advan- 
tageous to  substitute  either  calcium  or  a  mixture  of  calcium  oxide 
(lime),  magnesium,  and  sodium  for  the  magnesium  which  Ramsay  used 
for  removing  the  nitrogen.  The  commercial  oxygen  prepared  from 
liquid  air  (p.  18)  contains  as  much  as  3  per  cent  of  argon  and  serves 
as  the  most  economical  source  for  preparation  of  argon  in  quantities. 

Experiments  show  that  100  volumes  of  air  contain  about  0.94  vol- 
umes of  argon.  One  liter  of  argon  weighs  1.7809  g.,  so  that,  as  pre- 
dicted, it  is  heavier  than  nitrogen,  1 1.  of  which  weighs  but  1.2507  g. 
The  gas  is  characterized  by  its  complete  chemical  inertness.  Although 
subjected  to  the  action  of  many  other  substances  under  conditions 
which  have  been  found  most  favorable  for  effecting  the  combination 
of  elements,  it  has  not  as  yet  been  possible  to  prepare  any  compounds 
of  the  gas.  It  was  because  of  its  inertness  that  it  was  named  argon,  a 
word  meaning  "lazy"  or"  idle."  Like  nitrogen,  it  is  colorless,  odorless, 
and  tasteless.  It  has  been  condensed  to  a  liquid  which  boils  at  —  186°. 
At  still  lower  temperatures  it  forms  an  icelike  solid. 

Helium.  In  1889  the  American  chemist  Hillebrand  found  that 
certain  minerals  containing  the  element  uranium  evolve  a  gas  when 
heated.  He  concluded,  from  a  brief  investigation,  that  the  gas  so 
evolved  is  nitrogen.  In  1895,  shortly  after  the  discovery  of  argon, 
the  attention  of  Ramsay  was  called  to  this  experiment,  with  the  sug- 
gestion that  argon  might  be  present  in  the  gas  which  Hillebrand  had 
obtained.  Ramsay  repeated  the  experiment,  and  upon  examination  it 
was  found  that  the  spectrum  of  the  gas  contained  an  orange-colored 
line  identical  with  that  which  Lockyer,  in  1868,  had  detected  in  the 
spectrum  of  the  gases  surrounding  the  sun.  Lockyer  attributed  this 
line  to  the  presence  of  an  unknown  element  in  the  sun's  atmosphere, 
which  he  named  helium,  a  word  meaning  "  the  sun."  Ramsay's  experi- 
ments proved  that  the  gas  evolved  from  the  uranium  minerals  consist 
of  this  same  element,  helium,  mixed  with  small  .percentages  of  argon 
and  nitrogen.  Helium  has  since  been  found  in  the  gases  obtained  from 
many  minerals,  as  well  as  in  those  escaping  from  certain  springs.  Nat- 
ural gas  from  a  number  of  different  wells  has  also  been  examined  and 
helium  found  present  in  amounts  varying  from  traces  to  1.84  per  cent. 

Helium  is  commonly  avssociated  with  nitrogen  and  argon.  To  free 
it  from  nitrogen  the  gas  is  passed  over  heated  magnesium  or  calcium, 


110 


GENERAL  CHEMISTEY 


while  any  argon  present  is  separated  by  the  process  of  diffusion 
(see  p.  50).  Another  method  is  to  cool  the  gas  to  a  temperature 
of  about  —  200°,  at  which  temperature  both  argon  and  nitrogen  are 
liquefied  and  thus  separated  from  the  helium,  which  remains  in  the 
gaseous  state. 

Helium  resembles  argon  in  that  it  is  a  colorless,  odorless,  tasteless, 
and  inert  gas.  All  efforts  to  effect  a  combination  of  it  with  other  ele- 
ments have  failed.  With  the  exception  of  hydrogen,  it  is  the  lightest 
of  all  gases,  1  1.  of  it  weighing  0.1782  g.  Liquid  helium  boils  at 
—  268.7°  and  has  a  density  of  0.15.  It  may  be  added  that  of  all  known 
gases  helium  is  the  most  difficult  to  liquefy. 

Neon,  krypton,  and  xenon.  Following  the  discovery  of  argon  and 
helium  an  exhaustive  examination  of  various  gases  was  made,  espe- 
cially of  those  obtained  from  minerals,  in  the  hope  that  still  other 
elements  might  be  discovered.  These  investigations  proved  fruitless 
until  finally  directed  to  liquid  air  as  a  possible  source  of  such  unknown 
elements.  Large  quantities  of  liquid  air  were  subjected  to  careful  frac- 
tional distillation  and  the  different  fractions  examined,  especially  by 
the  spectroscope,  for  the  presence  of  unknown  elements.  Without 
going  into  detail  we  shall  simply  note  here  that  by  this  method  Ramsay 
and  Travers,  in  1898,  succeeded  in  isolating  three  new  elements, 
which  were  named  neon,  krypton,  and  xenon,  meaning,  respectively, 
"  new,"  "  hidden,"  and  "  stranger."  These  elements  proved  to  be  simi- 
lar to  helium  and  argon  in  being  entirely  devoid  of  chemical  activity. 

In  an  effort  to  discover  still  other  elements  in  air,  Moore  examined 
the  heavier  gases,  first  from  19  and  later  from  100  tons  of  liquid  air, 
but  found  no  new  elements. 

Some  of  the  main  facts  in  reference  to  these  inactive  gases  are 
given  in  the  following  table : 


HELIUM 

NEON 

ARGON 

KRYPTON 

XENON 

Weight  of  1  1.  (Watson)     .     . 

0.1782 

0.9002 

1.7809 

3.708 

5.851 

Number  of  cc.  dissolved  by  1  1. 

of  water  at  20°  (Antropoff) 

13.8 

14.7 

37.9 

73. 

110.9 

Boiling  point  of  liquid  form    . 

-  268.7° 

-239° 

-186° 

-  151.7° 

-109° 

Melting  point  of  solid  form     . 

>-270° 

-  188° 

-  169° 

-140° 

Number  of  volumes  in  1,000,- 

000  volumes  of  air  as  esti- 

mated by  Ramsay-     .     .     . 

4.00 

12.3 

9400 

0.05 

0.006 

NITROGEN  AND  THE  RARE  ELEMENTS 


111 


The  spectroscope.  The  spectroscope,  invented  by  Bunsen  and  Kir- 
choff  in  1860,  has  been  of  the  greatest  service  in  many  chemical 
investigations,  such  as  those  involved  in  the  discovery  and  isolation 
of  the  helium  group  of  gases,  so  that  a  brief  description  of  the  principle 
upon  which  the  instrument  is  constructed,  and  the  methods  employed 
in  its  use,  will  not  be  out  of  place  here. 

When  a  beam  of  light  passes  through  a  triangular  prism  of  glass,  it 
is  bent  out  of  its  course  and  emerges  at  a  decided  angle  with  its  original 

direction,  as  shown  in  Fig.  53.  Ordinary 
light  is  made  up  of  many  different  wave 
lengths,  and  each  one  is  deflected,  or  re- 
fracted, to  a  different  degree,  so  that  the 
various  colors  of  which  the  light  is  com- 
posed are  spread  out  in  a  series,  the  red 
being  the  least  refracted,  the  violet  the 

most  so.  A  beam  of  white  light  gives  a  continuous  series  of  colors 
from  red  through  orange,  yellow,  green,  blue,  to  violet,  called  a  con- 
tinuous spectrum.  The  spectrum  of  any  colored  light  is  not  continu- 
ous, but  shows  merely  those  colors  of  which  the  light  is  composed. 

That  these  colors  may  be  made  as  distinct  and  sharply  separated  as 
possible,  the  light  should  shine  upon  the  pri§m  through  a  very  narrow 
slit  in  a  screen,  ar- 
ranged so  as  to  be  par- 
allel with  the  axis  of 
the  prism.  The  colors 
will  then  be  a  series  of 
narrow  lines,  each 


an 


image  of  the  slit,  spread 

out  parallel  with  each 

other.    An  instrument, 

the   essential   parts  of 

which  are  a   prism,  a 

screen  provided  with  a 

narrow  slit,  and  lenses 

for  focusing  the  light 

upon  the  slit   and  for 

viewing  the  spectrum,  is  called  a  spectroscope,  or  spectrometer.  Fig.  54 

represents  a  simple  form  of  such  a  spectroscope,  the  slit  being  seen  at 

the  end  of  the  tube  B.  When  we  look  into  the  eyepiece  A,  the  spectrum 


FIG.  54 


112 


GENERAL   CHEMISTEY 


of  the  flame  is  seen  as  a  series  of  bright  lines  on  a  dark  field.  The 
tube  C  contains  a  scale  which  is  also  seen  when  we  look  into  the  eye- 
piece. Any  incandescent  solid,  such  as  a  glowing  platinum  wire,  glow- 
ing carbon,  or  an  incandescent  lamp,  gives  a  continuous  spectrum,  all 
wave  lengths  of  light  being  represented.  Light  from  volatilized  salts 
and  glowing  gases  gives  an  interrupted  or  line  spectrum,  characteristic 
of  the  particular  substance  giving  rise  to  the  light.  Each  element  and 
many  stable  compounds  may  therefore  be  recognized  by  their  spectrum. 
Methods  of  use.  There  are  a  number  of  ways  in  which  these  prin- 
ciples may  be  applied  in  practice,  depending  on  the  way  in  which 
the  substance  under  investigation  is  brought  to  the  point  of  giving 
out  light. 

1.  Flame  spectrum.    The  simplest  way  is  to  dip  a  platinum  wire 
into  some  of  the  substance,  or  a  concentrated  solution  of  it,  and  heat 
the  wire  in  a  Bunsen  flame.  The  flame  takes  on  the  color  of  the  glow- 
ing vapor  of  the   substance,   and  gives   its  spectrum  when  viewed 
through  the  spectroscope.    When  a  higher  temperature  is  desired,  a 
little  cavity  may  be  hollowed  out  in  the  lower  carbon  of  an  arc  light, 
and  some  of  the  solid  placed  in  this,  the  heat  of  the   arc  slowly 
volatilizing  the  solid. 

2.  Vacuum-tube  spectruhi.  When  a  gas  is  sealed  within  a  tube  pro- 
vided with  electrodes  (known  as  a  Pliicker  tube),  such  as  is  repre- 
sented in  A  (Fig.  55),  the  pressure  of 
the  gas  having  foeen  reduced  to  only 
a  few  millimeters  of  mercury  before 
the  tube  was  sealed,  a  high  voltage 
applied  to  the  electrodes  will  cause 
the  rarefied  gas  to  become  brilliantly 
luminous.    The  light  in  the  capillary 
portion  of  the  tube  can  be  viewed 
to   advantage    in    the    spectroscope, 
and  a  very  characteristic  spectrum  is 
found  for  each  gas.  The  character  of 
each  spectrum  is  much  influenced  by 
the  electrical  conditions  and  by  the 

pressure  of  the  gas,  and  the  variations  so  produced  assist  in  the  iden- 
tification of  any  particular  gaseous  substance.  Argon,  for  example, 
gives  a  very  complex  series  of  lines..  With  an  intermittent  discharge 
the  glow  in  the  tube  is  red  and  the  red  lines  in  the  spectrum  are  very 


FIG.  55 


NITROGEN  AND  THE  RARE  ELEMENTS  113 

brilliant;  with  an  oscillating  discharge  the  glow  is  bright  blue,  the 
red  spectrum  lines  largely  disappear,  and  many  new  green  and  blue 
ones  come  to  view.  With  helium  under  a  pressure  of  from  7  to  8  mm. 
the  glow  is  bright  yellow ;  if  the  pressure  is  reduced  to  from  1  to 
2  mm.,  the  tube  emits  a  green  light.  The  current  is  usually  supplied 
by  an  induction  coil  B,  operated  by  a  battery  (Fig.  55). 

3.  Absorption  spectrum.  It  has  been  stated  that  an  incandescent 
solid  gives  a  continuous  spectrum.  If  a  layer  of  gas  is  interposed  be- 
tween the  incandescent  solid  and  the  prism,  the  gas  will  absorb  those 
wave  lengths  of  light  which  it  can  itself  give  out  when  luminous,  and 
the  continuous  spectrum  will  be  interrupted  by  a  series  (3f  black  lines 
corresponding  to  the  bright  lines  of  the  gas.  In  this  way  the  solar 
spectrum  is  crossed  by  a  great  number  of  dark  lines  due  to  the  pres- 
ence of  gases  in  the  envelope  surrounding  the  incandescent  center  of 
the  sun.  Such  lines  are  called  Fraunhofer  lines,  after  then*  discoverer, 
and  a  spectrum  in  which  they  appear  is  called  an  absorption  spectrum. 
We  have  learned  much  about  the  elements  present  in  the  sun  from  a 
study  of  its  absorption  spectrum. 


CHAPTER  X 

THE  ATMOSPHERE 

Atmosphere  and  air.  The  term  atmosphere  is  applied  to  the  gaseous 
envelope  surrounding  the  earth;  the  term  air  is  generally  applied 
to  a  limited -portion  of  this  envelope,  although  the  two  terms  are 
often  used  interchangeably.  Many  references  have  been  made  to  the 
composition  and  properties  of  the  atmosphere.  These  statements  will 
now  be  collected  and  discussed  somewhat  more  in  detail. 

Historical.  Like  water,  air  was  first  regarded  as  elementary  in  char- 
acter. Near  the  close  of  the  eighteenth  century  Scheele,  Priestley, 
Cavendish,  and  Lavoisier,  by  their  experiments,  showed  it  to  be  a 
mixture  of  at  least  two  gases  —  those  which  we  now  call  oxygen  and 
nitrogen.  By  absorbing  the  oxygen  from  an  inclosed  volume  of  air, 
and  noting  the  contraction  in  volume  due  to  the  removal  of  oxygen, 
Cavendish  was  able  to  determine  with  considerable  accuracy  the 
relative  volumes  of  oxygen  and  nitrogen  present. 

The  constituents  of  the  air.  The  normal  constituents  of  air,  to- 
gether with  the  approximate  amounts  of  each  in  samples  collected  in 
the  open  fields,  are  as  follows : 

Oxygen 21  volumes  in  100  volumes  of  dry  air 

Nitrogen 78  volumes  in  100  volumes  of  dry  air 

Water  vapor variable  within  wide  limits 

Carbon  dioxide  .     .     .     .  •  .     .     .  3  to  4  volumes  in  10,000  volumes  of  dry  air 

Argon   .  0.937  volumes  in  100  volumes  of  dry  air 

Helium,  neon,  krypton,  xenon  .     .  traces 

There  are  also  present  small  quantities  of  hydrogen  peroxide,  am- 
monium nitrate,  dust  particles  of  various  kinds,  microorganisms, 
and  probably  traces  of  hydrogen  and  ozone.  In  addition  to  these 
constituents  the  air  in  large  cities  and  manufacturing  districts 
contains  varying  amounts  of  substances  such  as  hydrogen  sulfide 
(H2S),  sulfur  dioxide  (SO2),  and  carbon  monoxide  (CO),  which 
are  evolved  in  the  decay  of  matter  or  are  formed  in  certain  manu- 
facturing processes. 

114 


THE  -  ATMOSPHERE  115 

For  the  purposes  of  discussion  it  is  convenient  to  divide  the  con- 
stituents of  the  air  into  two  general  classes,  the  one  including  those 
essential  to  life  and  the  other  those  not  essential. 

The  essential  constituents.  The  constituents  which  are  known  to  be 
essential  to  life  are  oxygen,  nitrogen,  water  vapor,  and  carbon  dioxide. 
The  first  three  of  these  have  already  been  discussed  in  detail.  The 
remaining  one,  carbon  dioxide,  often  called  carbonic  acid  gas,  is  a  gas 
having  the  formula  CO2.  Reference  has  already  been  made  to  the 
fact  that  it  is  evolved  in  the  processes  of  both  respiration  and  com- 
bustion, so  that  large  quantities  of  it  are  continually  being  added  to 
the  atmosphere.  The  properties  of  the  gas  will  be  described  in  the 
chapter  relating  to  the  compounds  of  carbon ;  it  is  only  necessary  to 
note  here  that  it  is  a  comparatively  heavy  gas  and  will  neither  burn 
nor  support  combustion. 

The  oxygen  in  the  atmosphere  directly  supports  life  through  the 
process  of  respiration.  The  nitrogen  serves  to  dilute  the  oxygen  and 
thus  to  diminish  the  intensity  of  its  action.  It  is  likewise  assimilated 
by  certain  plants  (see  p.  107).  The  water  vapor  prevents  excessive 
evaporation  of  the  water  present  in  organisms,  while  the  carbon  diox- 
ide is  an  essential  plant  food. 

The  nonessential  constituents.  It  is  unnecessary  to  enter  into  any 
extended  discussion  of  these  constituents.  Some  of  them,  such  as 
the  inactive  gases  described  in  the  preceding  chapter,  are  always 
present,  while  others,  such  as  sulfur  dioxide  and  hydrogen  sulfide, 
are  found  in  certain  localities  only  and  may  be  classed  as  accidental 
constituents. 

Hydrogen  peroxide  is  formed  in  minute  quantities  in  certain  proc- 
esses of  oxidation  which  are  constantly  taking  place,  and  the  traces  of 
it  in  the  atmosphere  are  probably  derived  from  this  source.  The  am- 
monium nitrate  present  is  formed  by  the  action  of  ammonia  (NH3) 
upon  nitric  acid  (HNO3).  The  former  of  these  is  a  gas  evolved  in  the 
decay  of  nitrogenous  matter;  the  latter  results  from  the  action  of 
water  vapor  on  an  oxide  of  nitrogen  formed  by  the  combination  of 
oxygen  and  nitrogen  through  the  effect  of  lightning  discharges  (see 
p.  107).  The  composition  of  the  dust  particles  differs  greatly,  since 
almost  any  sort  of  finely  divided  material  will  float  in  the  air.  It  is 
interesting  to  note  that  these  particles  play  a  conspicuous  part  in  the 
formation  of  dew  and  rain,  since  they  serve  as  centers,  or  nuclei,  for 
the  condensation  of  moisture. 


116 


GENERAL   CHEMISTEY 


Just  as  microorganisms  occur  in  all  natural  waters,  so  we  always 
find  them  present  in  the  air.  While  it  is  undoubtedly  true  that  trans- 
missible diseases  are  sometimes  spread  by  these  organisms,  it  must 
not  be  inferred  that  their  presence  is  always  a  menace  to  health.  The 
processes  of  putrefaction  and  fermentation  are  due  to  these  organisms 
and  may  be  prevented  by  excluding  them.  Thus,  in  the  ordinary 
processes  of  canning  fruits  the  materials  are  first  heated  and  then 
sealed  in  air-tight  receptacles.  The  heat  kills  any  organisms  already 
present,  while  the  exclusion  of  air  prevents  their  subsequent  intro- 
duction from  that  source.  Preservatives  are  simply  substances  that 
prevent  the  growth  of  these  organisms. 

Methods  used  for  determining  the  percentages  of  oxygen,  nitrogen, 
water  vapor,  and  carbon  dioxide  present  in  the  atmosphere.  The  first 
determinations  of  the  relative  amounts  of  oxygen  and  nitrogen  in  air 
were  made  by  Cavendish  in  1783.  The  method  employed  consisted 
in  the  removal  of  oxygen  from  a  definite  volume  of  air  by  means  of 
nitric  oxide,  a  compound  which  possesses  the  property  of  combining 
with  oxygen  to  form  a  gas  known  as  nitrogen  dioxide,  which  is  absorbed 
by  water.  The  volume  of  oxygen  present  in  the  air  was  computed 
from  the  contraction  in  the  volume  of  the  air  due  to  its  removal. 
This  method,  however,  has  given  way  to  other  more  accurate  ones,  the 
most  important  of  which  are  the  following : 

1.  Determination  of  oxygen.  The  general  method  used  consists  in 
removing  the  oxygen  from  a  definite  volume  of  air  and  noting  the 
contraction  in  volume.  Phosphorus  is  gener- 
ally used  for  removing  the  oxygen,  although 
copper  and  hydrogen  are  sometimes  employed. 
If  great  accuracy  is  required,  the  carbon  di- 
oxide present  in  the  sample  of  air  is  first 
removed  by  a  suitable  reagent,  the  com- 
pound known  as  potassium  hydroxide  being 
the  one  generally  used. 

When  phosphorus  is  used  as  an  absorbent  for  the 
oxygen,  the  determination  may   be   carried   out  as 
follows :  A  tube  is  filled  with  water  and  inverted 
-p       r(>  in  a  vessel  of    water    (Fig.  56).    A    sample  of  the 

air  to  be  analyzed  is  then  introduced  into  the  tube 

until  it  is  partly  filled.  The  volume  of  the  inclosed  air  is  carefully  noted  and 
reduced  to  standard  conditions.  A  piece  of  phosphorus  is  attached  to  a  wire  and 
brought  within  the  tube,  as  shown  in  the  figure.  The  oxygen  combines  with  the 


THE ,  ATMOSPHERE  11T 

phosphorus,  water  rising  to  take  the  place  of  the  absorbed  oxygen.  The  com- 
plete removal  of  the  oxygen  is  indicated  by  the  fact  that  no  further  diminu- 
tion in  the  volume  of  the  gas  in  the  tube  takes  place.  The  phosphorus  is  then 
removed  and  the  volume  of  the  residual  gas  determined  and  reduced  to  standard 
conditions.  The  contraction  in  the  volume  of  the  air  is  equal  to  the  volume  of 
oxygen  withdrawn.  The  time  required  for  the  removal  of  the  oxygen  is  greatly  re- 
duced in  practice  by  using  a  number  of  strips  or  wires  of  phosphorus  (see  p.  28), 
as  well  as  by  using  a  form  of  apparatus  which  makes  it  possible  to  cause  the 
inclosed  air  to  flow  back  and  forth  over  the  phosphorus. 

A  more  accurate  method  is  the  following,  in  which  the  oxygen  is 
removed  by  combination  with  hydrogen.  A  eudiometer  tube  is  filled 
with  -mercury  and  inverted  in  a  vessel  of  the  same  liquid.  A  conven- 
ient amount  of  air  is  then  introduced  into  the  tube  and  its  volume 
accurately  noted.  There  is  then  introduced  more  than  sufficient  hy- 
drogen to  combine  with  the  oxygen  present  in  the  inclosed  air,  and 
the  volume  is  again  accurately  noted.  The  mixture  is  then  exploded 
by  an  electric  spark,  and  the  volume  is  once  more  taken.  By  sub- 
tracting this  volume  from  the  total  volume  of  the  air  and  hydrogen 
there  is  obtained  the  contraction  in  volume  due  to  the  union  of  the 
oxygen  and  hydrogen.  The  volume  occupied  by  the  water  formed  by 
the  union  of  the  two  gases  is  so  small  that  it  may  be  disregarded  in 
the  calculation.  Since  oxygen  and  hydrogen  combine  in  the  ratio  1  to 
2  by  volume,  it  is  evident  that  one  third  of  the  total  contraction  is 
equal  to  the  volume  originally  occupied  by  the  oxygen  in  the  inclosed 
air.  The  following  example  will  make  this  clear : 

Volume  of  air  in  tube .     .  50.0  cc. 

Volume  after  introducing  hydrogen 80.0  cc. 

Volume  after  combination  of  oxygen  and  hydrogen 48.5  cc. 

Contraction  in  volume  due  to  combination  ....  (80  cc.  —  48.5  cc.)  31.5  cc. 
Volume  of  oxygen  in  50  cc.  of  air  (^  of  31.5) 10.5  cc. 

All  these  methods  agree  in  showing  that  100  volumes  of  dry  air 
contain  21  volumes  of  oxygen,  with  a  variation  of  not  more  than 
0.2  per  cent. 

2.  Determination  of  nitrogen.  If  the  gas  left  after  the  removal  of 
oxygen  from  a  sample  of  air  is  passed  over  heated  magnesium,  the 
nitrogen  is  withdrawn,  argon  and  the  other  rare  elements  being  left. 
It  may  thus  be  shown  that  of  the  79  volumes  of  gas  le|f  after  the  re- 
moval of  the  oxygen  from  100  volumes  of  air,  approximately  78  are 
nitrogen  and  1.0  argon.  The  other  elements  are  present  in  such  small 
quantities  that  they  may  be  neglected. 


118  GENERAL  CHEMISTEY 

3.  Determination  of  carbon  dioxide.    The  method  ordinarily  used  in 
the  determination  of  carbon  dioxide  consists  in  shaking  a  definite 
volume  of  air  with  a  solution  of  barium  hydroxide  (Ba(OH)2).    The 
carbon  dioxide  reacts  with  the  barium  hydroxide  to  form  barium  car- 
bonate (BaCO3),  which  is  insoluble  in  water,  so  that  it  may  be  filtered 
off  and  its  weight  determined.    From  this  weight  it  is  easily  possible 
to  calculate  the  percentage  of  the  carbon  dioxide  which  was  present 
in  the  air. 

4.  Determination  of  water  vapor.    The  water  vapor  present  in  a  given 
volume  of  air  may  be  determined  by  passing  the  air  over  calcium  chlo- 
ride (or  some  other  compound  which  readily  absorbs  moisture)  and  not- 
ing the  increase  in  the  weight  of  the  chloride.    The  percentage  present 
not  only  varies  with  the  locality,  but  varies  widely  from  day  to  day  in 
the  same  locality,  because  of  the  winds  and  changes  in  temperature. 

Processes  affecting  the  composition  of  the  air.  The  most  important 
of  these  processes  are  the  following : 

1.  Respiration.    In  the  process  of  respiration  some  of  the  oxygen  in 
the  inhaled  air  is  absorbed  by  the  blood  and  carried  to  all  parts  of  the 
body,  where  it  combines  with  the  hydrogen  and  carbon  of  the  worn-out 
tissues.    The  products  of  oxidation  are  carried  back  to  the  lungs  and 
exhaled  in  the  form  of  moisture  and  carbon  dioxide.   The  volume  of 
carbon  dioxide  exhaled  by  an  adult  averages  about  20  1.  per  hour. 
Hence,  in  a  poorly  ventilated  room  occupied  by  a  number  of  people 
the  percentage  of  carbon  dioxide  rapidly  increases.    While  this  gas 
is  not  poisonous  unless  present  in  large  amounts,  nevertheless  air 
containing  more  than  15  parts  in  10,000  is  not  fit  for  respiration. 

2.  Combustion.    All  the  ordinary  forms  of  fuel  contain  large  percen- 
tages of  carbon.   On  burning,  this  carbon  combines  with  the  oxygen  in 
the  air,  forming  carbon  dioxide. 

3.  Decay  of  organic  matter.    The  decay  of  organic  matter  is  largely 
a  process  of  oxidation  brought  about  through  the  agency  of  micro- 
organisms.   The  carbon  and  hydrogen  are  evolved  principally  in  the 
form  of  carbon  dioxide   and  water.    Any  nitrogen  present  may  be 
evolved  as  free  nitrogen,  ammonia,  or  various  oxides  of  nitrogen, 
depending  upon  the  conditions.    The  effect  of  the  decay  of  organic 
matter  upon  the  composition  of  the  atmosphere  is  therefore  similar 
to  respiration  and  combustion  in  that  it  tends  to  diminish  the  per- 
centage of  free  oxygen  and  to  increase  correspondingly  that  of  carbon 
dioxide  .present. 


THE  ATMOSPHERE 


119 


4.  Action  of  plants.    Plants  have  the  power,  when  in  the  sunlight, 
of  absorbing  carbon  dioxide  from  the  air,  retaining  the  carbon  and  re- 
turning a  portion  of  the  oxygen  to  the  air.    It  will  be  observed  that 
these  changes  are  just  the  opposite  of  those  brought  about  by  the 
processes  of  respiration,  combustion,  and  decay. 

5.  Weathering  of  rocks.   Large  quantities  of  carbon  dioxide  are  being 
constantly  withdrawn  from  the  atmosphere  through  its  combination 
with  various  rock  materials. 

The  constancy  of  the  composition  of  air.  Notwithstanding  the  changes 
constantly  taking  place  which  tend  to  alter  the  composition  of  the  air, 
the  results  of  a  great  many  analyses  of  air  collected  in  the  open  fields 
show  that  the  percentages  of  oxygen  and  nitrogen,  as  well  as  of  carbon 
dioxide,  are  very  nearly  constant.  This  constancy  of  composition  in 
the  case  of  oxygen  is  shown  in  the  following  analyses  tabulated  by 
Clarke.  The  percentages  are  expressed  in  volumes. 


LOCALITY  OF  SAMPLES 

NUMBER  OF 
ANALYSES 

MINIMUM 

MAXIMUM 

Paris 

100 

20.913 

20.999 

Heidelberg     

28 

20.840 

20.970 

Manchester    

32 

20.780 

21.020 

Mountains  of  Scotland 
Dresden     
Cape  Horn 

34 
46 

20 

20.800 

20.877 
20.72 

21.180 
20.971 
20.970 

Cleveland,  Ohio       .... 

45 

20.900 

20.950 

Since  the  percentages  of  oxygen  and  nitrogen  in  the  air  are  so  con- 
stant, the  question  naturally  arises,  whether  these  two  elements  are 
not  present  in  the  form  of  a  definite  chemical  compound.  That  they 
are  not  combined,  but  are  simply  mixed  together,  can  be  shown  in  a 
number  of  ways,  among  which  are  the  following : 

1.  When  air  dissolves  in  water,  it  has  been  found  that  the  ratio  of 
oxygen  to  nitrogen  in  the  dissolved  air  is  no  longer  21 :  78  but  more 
nearly  35 :  65.    If  it  were  a  chemical  compound,  the  ratio  of  oxygen 
to  nitrogen  would  not  be  changed  by  solution. 

2.  A  chemical  compound  in  the  form  of  a  liquid  has  a  definite 
boiling  point.   Water,  for  example,  boils  at  100°.   Moreover,  the  steam 
which  is  thus  formed  has  the  same  composition  as  the  water.    The 
boiling  point  of  liquid  air,  on  the  other  hand,  gradually  rises  as  the 
liquid  boils,  the  nitrogen  escaping  first,  followed  by  the  oxygen.    If 


120  GENERAL  CHEMISTRY 

the  two  were  combined,  they  would  pass  off  together  in  the  ratio  in 
which  they  are  found  in  the  air. 

Why  the  air  has  a  constant  composition.  If  air  is  a  mixture  and 
changes  are  constantly  taking  place  which  tend  to  modify  its  compo- 
sition, how,  then,  do  we  account  for  the  constancy  of  composition 
which  the  analyses  reveal  ?  This  is  explained  by  several  facts : 
(1)  the  changes  which  are  caused  by  the  processes  of  combustion, 
respiration,  and  decay,  on  the  one  hand,  and  the  action  of  plants,  on 
the  other,  tend  to  equalize  each  other;  (2)  the  winds  keep  the  air 
in  constant  motion  and  so  prevent  local  changes ;  (3)  the  total 
quantity  of  oxygen  and  carbon  dioxide  in  the  atmosphere  is  so  large 
as  to  be  practically  unaffected  by  any  of  the  naturally  occurring 
changes. 

Impure  air.  Priestley,  in  his  experiments  on  air,  observed  that  when 
nitric  oxide  was  introduced  into  air  confined  over  water,  a  contraction 
in  the  volume  of  the  air  took  place  which  he  correctly  attributed  to 
the  absorption  of  the  "  dephlogisticated  air "  (oxygen)  by  the  nitric 
oxide.  This  method  was  therefore  used  to  determine  the  purity  of 
air  (or  the  "  goodness,"  as  Priestley  termed  it),  upon  the  assumption 
that  the  percentage  of  oxygen  present  in  any  sample  of  air  is  a  direct 
measure  of  its  purity.  The  kind  of  tube  in  which  the  air  was  confined 
in  making  this  determination  was  termed  a  eudiometer  —  a  word 
meaning  "  a  measure  of  goodness,"  and  this  term  is  still  retained. 

It  is  now  known  that  air  becomes  impure  or  foul  through  breathing, 
not  because  of  lack  of  oxygen,  nor  ordinarily  because  of  the  presence 
of  an  abnormal  amount  of  carbon  dioxide,  but  owing  rather  to  the 
organic  matter  thrown  off  from  the  lungs  in  the  exhaled  air.  Never- 
theless, the  purity  of  such  air  is  commonly  judged  by  the  amount  of 
carbon  dioxide  present,  for  the  percentage  of  this  gas  is  easily  deter- 
mined and  serves  as  an  indirect  measure  of  the  amount  of  organic 
matter  present,  since  the  two  are  exhaled  together. 

The  approximate  differences  in  the  percentages  of  oxygen,  carbon 
dioxide,  and  water  vapor  in  inhaled  and  exhaled  air  are  shown  in 
the  following  analyses: 


INHALED  AIR 

EXHALED  AIR 

Oxygen    , 

21.00% 

16.00% 

Carbon  dioxide       ....          .     . 

0.04% 

4.38% 

Moisture       

variable 

saturated 

THE  ATMOSPHERE  121 

The  cycle  of  carbon  in  nature.  Under  the  influence  of  sunlight  the 
carbon  dioxide  absorbed  from  the  air  by  plants  reacts  with  water  and 
small  amounts  of  other  substances  absorbed  from  the  soil  to  form 
complex  compounds  of  carbon  which  constitute  the  essential  part  of 
the  plant  tissue.  This  reaction  is  attended  by  the  evolution  of  oxygen, 
which  is  restored  to  the  air.  The  compounds  resulting  from  these 
changes  are  much  richer  in  their  energy  content  than  are  the  sub- 
stances from  which  they  are  formed ;  hence,  a  certain  amount  of 
energy  must  have  been  absorbed  in  their  formation.  The  source  of 
this  energy  is  the  sun's  rays. 

If  the  plant  is  burned  or  decays  in  the  open  ah",  the  changes  which 
took  place  in  the  formation  of  the  compounds  present  are  largely 
reversed.  The  carbon  and  hydrogen  combine  with  oxygen  taken  from 
the  air  to  form  carbon  dioxide  and  water,  while  the  energy  absorbed 
from  the  sun's  rays  is  liberated  in  the  form  of  heat  energy.  If,  on 
the  other  hand,  the  plant  is  used  as  food,  the  compounds  present  are 
utilized  in  building  up  the  tissues  of  the  body  and  as  a  source  of 
energy.  In  either  case  the  carbon  and  hydrogen  ultimately  combine 
with  inhaled  oxygen  to  form  carbon  dioxide  and  water,  which  are  in 
turn  exhaled.  The  energy  possessed  by  the  food  substance  is  lib- 
erated partly  in  the  form  of  heat,  which  maintains  the  temperature 
of  the  body,  and  partly  as  muscular  energy.  The  carbon  originally 
absorbed  from  the  air  by  the  plant  in  the  form  of  carbon  dioxide 
is  thus  restored,  and  the  cycle  of  changes  may  be  repeated. 

The  properties  of  air.  Inasmuch  as  air  is  composed  principally  of  a 
mixture  of  oxygen  and  nitrogen,  which  elements  have  been  discussed 
in  detail,  its  properties  may  be  inferred  largely  from  those  of  the  two 
gases.  The  preparation  and  properties  of  liquid  air  and  its  use  as  a 
source  of  oxygen,  nitrogen,  and  the  rare  elements  in  the  atmosphere 
have  likewise  been  discussed. 


CHAPTER  XI 

SOLUTIONS 

Introduction.  In  Chapter  I  a  distinction  was  made  between  a  mixture 
and  a  compound.  In  a  typical  mixture  particles  of  different  properties 
may  be  distinguished,  so  it  is  not  of  perfectly  uniform  character.  In 
a  compound  every  smallest  portion  is  identical  in  composition  with 
every  other  portion. 

Intermediate  between  these  is  a  great  class  of  bodies  called  solutions. 
the  most  familiar  types  of  which  are  solutions  of  solids  in  liquids. 
They, differ  most  noticeably  from  mixtures  in  that  they  are  of  per- 
fectly even  character  throughout,  which  fact  is  usually  expressed  by 
saying  that  they  are  homogeneous.  They  differ  from  definite  chemical 
compounds  in  that  their  composition  can  be  varied  between  wide  lim- 
its. A  solution  may  therefore  be  denned  as  a  body  of  homogeneous 
character  whose  composition  may  be  varied  continuously  between  cer- 
tain limits.  This  definition  makes  no  restrictions  as  to  the  physical 
state  of  the  solution  or  of  its  constituents.  It  includes  any  combi- 
nation, such  as  gases  in  gases  or  in  liquids,  solids  in  liquids  or  in 
solids.  Since  there  are  three  general  states  of  matter,  there  are  evi- 
dently nine  possible  combinations,  some  of  which  are  of  small  impor- 
tance. It  will  be  sufficient  for  our  purposes  to  consider  five  of  the 
most  important  pairs. 

Solutions  of  gases  in  gases.  We  rarely  think  of  a  mixture  of  two 
gases  as  a  solution,  yet  it  conforms  to  our  definition  and  in  many  ways 
recalls  the  more  familiar  case  of  the  solution  of  one  liquid  in  another. 
Regarding  it  as  a  solution,  we  find  its  properties  to  be  of  a  very  sim- 
ple character.  All  gases  mix  in  every  proportion  and  apparently  have 
no  effect  upon  each  other  save  when  obvious  chemical  action  occurs. 
If  their  volumes  and  pressures  are  equal  before  mixing,  the  mixture 
will  have  the  double  volume  if  the  pressure  remains  constant.  Of  this 
pressure  one  half  will  be  due  to  each  gas,  so  that  each  will  exert  the 
same  pressure  as  if  it  alone  had  been  put  in  the  double  volume  and 
the  other  were  not  present.  When  dealing  with  a  solution  of  gases, 
therefore,  we  may  assign  to  each  its  fraction  of  the  total  pressure,  which 

122 


SOLUTIONS  123 

will  be  the  same  as  though  the  other  gases  were  to  be  removed,  leav- 
ing the  one  alone  in  the  inclosing  volume.  This  fraction  of  the  pres- 
sure is  called  the  partial  pressure  of  the  gas.  Thus,  as  we  have  seen, 
the  partial  pressure  of  the  aqueous  vapor  in  air  standing  over  water 
at  20°  is  17.4  mm.  Since  there  is  no  expansion  or  contraction  when 
gases  are  mixed,  the  density  of  the  mixture  can  be  calculated  from  the 
known  densities  and  percentages  of  the  constituent  gases. 

Solutions  of  gases  in  liquids.  In  this  and  in  the  succeeding  types  of 
solutions,  one  of  the  constituents  acts  as  the  dissolving  body  and  is 
termed  the  solvent;  the  other  acts  as  the  dissolved  body,  or  solute. 
While  gases  mix  with  each  other  in  all  proportions,  a  liquid  will  take 
up  but  a  limited  quantity  of  a  gas.  When  the  gas  has  been  bubbled 
through  the  liquid  until  no  more  is  dissolved,  the  liquid  is  said  to  be 
saturated.  The  ratio  of  the  quantity  of  the  gas  solute  to  the  liquid 
solvent  is  called  the  solubility  of  the  gas.  This  may  be  expressed  in  a 
variety  of  ways,  the  most  prevalent  usage  being  to  state  the  number 
of  volumes  of  gas  (under  definite  conditions  of  temperature  and 
pressure)  dissolved  in  one  volume  of  the  solvent.  The  table  shows 
the  solubility  in  water  of  a  few  of  the  more  familiar  gases. 

SOLUBILITY  OF  GASES  IN  WATER 


NAME  OF  GAS 

VOLUME  ABSORBED  AT  0°  AND 
UNDER  760  MM.  PRESSURE 

Ammonia 

1298  9  1 

Hydrogen  chloride 

506  0  1 

79.79  1. 

4.37  1. 

Carbon  dioxide      . 

1.713  1. 

Oxvsren 

0  0496  1. 

Hydrogen 

0  0214  1 

Nitrogen 

•  0  0233  1 

In  respect  to  solubility,  gases  fall  roughly  into  two  classes:  those 
of  rather  small  solubility,  such  as  oxygen,  hydrogen,  and  nitrogen; 
those  of  quite  large  solubility,  such  as  ammonia,  hydrogen  chloride, 
and  sulfur  dioxide.  These  groups  will  be  discussed  separately. 

Conditions  affecting  the  solubility  of  moderately  soluble  gases.  A 
number  of  different  factors  affect  the  solubility  of  a  moderately 
soluble  gas. 


124 


GENEKAL   CHEMISTRY 


1.  Effect  of  specific  properties.    The  specific  properties  of  both  gas 
and  solvent  are  of  the  first  importance  in  determining  the  solubility  of 
a  given  gas  in  a  given  solvent.    All  gases  are  soluble  to  some  extent 
in  every  solvent.    Other  conditions  being  equal,  no  two  gases  have  the 
same  solubility  in  a  given  liquid. 

2.  Effect  of  pressure;  the  law  of  Henry.    Increase  of  pressure  always 
increases  the  weight  of  gas  going  into  solution,  the  increase  being 
proportional  to  the  pressure.    This  is  known  as  the  law  of  Henry, 
having  been  formulated  by  him  in  1803.    If  1  g.  of  a  gas  dissolves 
in  100  cc.  of  water  at  atmospheric  pressure,  2  g.  will  dissolve  under 
2  atmospheres,  provided  the  temperature  remains  constant. 

3.  Effect  of  temperature.    With  rise  in  temperature,  gases  become  less 
soluble,  and  at  the  boiling  point  of  the  solvent  most  gases  are  insol- 
uble.   This  is  probably  not  due  to  genuine  insolubility  but  to  the  cir- 
cumstance explained  a  little   later.     The  solubility  of  a  gas  is  not 
exactly  inversely  proportional  to  the  temperature,  so  its  value  cannot 
be  calculated,  but  must    be  determined  experimentally  at  frequent 

temperature  intervals. 
The  results  of  such 
experiments  are  best 
tabulated  in  the  form 
of  curves  (Fig.  57). 

Temperature  is  repre- 
sented on  the  abscissa,  and 
concentrations,  in  grams 
per  liter,  on  the  ordinate. 
All  the  curves  show  dif- 
ferent solubilities  at  0°, 
which  diminish  as  the 

temperature  rises.    If  solubility  were  accurately  proportional  to  temperature,  it 
would  be  represented  by  a  straight  line  instead  of  by  a  curve. 

Conditions  affecting  the  solubility  of  very  soluble  gases.  In  a  quali- 
tative way  very  soluble  gases  are  affected  by  temperature  and 
pressure  in  the  same  way  as  are  sparingly  soluble  ones,  but  quanti- 
tatively the  law  of  Henry  does  not  hold,  the  solubility  not  being  pro- 
portional to  the  pressure.  The  specific  properties  of  both  gas  and 
solvent  play  a  more  prominent  part,  leading  to  very  wide  differences 
in  solubility,  as  may  be  seen  in  the  table  on  page  123.  There  is  always 
a  very  considerable  change  in  the  volume  of  the  solvent  when  a  large 
volume  of  gas  is  absorbed,  so  it  is  not  possible  to  calculate  the  density 


100 


50 


20°     30 


40°    50*    60° 
FIG.  57 


70"    60°    90"    100"  T 


SOLUTIONS  125 

which  such  a  solution  will  have,  nor  from  its  measured  density  to 
infer  the  weight  of  gas  it  has  absorbed.  Thus,  1  1.  of  water,  when 
saturated  with  ammonia  at  14°,  increases  in  volume  to  1.580  1.,  while 
its  density  decreases  to  0.8844. 

Solutions  of  this  kind  frequently  acquire  properties  possessed  by 
neither  the  gas  nor  the  liquid.  Thus,  neither  ammonia  nor  water  has 
any  effect  on  the  color  of  red  litmus  (an  organic  dye),  whereas  the 
solution  turns  it  blue  ;  similarly,  sulfur  dioxide  has  no  effect  on  blue 
litmus,  while  its  solution  turns  it  red.  For  these  and  many  other 
reasons  it  appears  probable  that  a  chemical  action  takes  place  between 
the  gas  and  the  solvent,  which  accounts  for  such  high  solubility.  On 
the  other  hand,  when  the  two  solutions  just  mentioned  are  exposed 
to  the  open  air,  ammonia  and  sulfur  dioxide  respectively  escape  from 
the  solution,  just  as  do  gases  which  give  no  evidence  of  chemical 
combination. 

Chemical  equilibrium  in  solution.  The  apparently  opposite  conclu- 
sions reached  in  the  last  paragraph  may  be  reconciled  by  assuming 
that  the  conditions  existing  in  the  liquid  resulting  from  passing 
ammonia  into  water  are  represented  by  the  equilibrium  equation 


which  indicates  that  the  ammonia  combines  with  water  to  form  a 
compound,  ammonium  hydroxide  (NH4OH),  but  that  the  reaction  is 
incomplete,  owing  to  the  decomposition  of  the  compound  into  its  orig- 
inal components.  The  equation  therefore  represents  an  equilibrium 
similar  to  that  between  oxygen  and  ozone,  or  between  water  and  its 
constituents,  oxygen  and  hydrogen.  In  a  subsequent  chapter  the  con- 
ditions which  result  in  such  types  of  equilibrium  will  be  discussed 
more  at  length. 

Solubility  of  mixed  gases  ;  the  law  of  Dalton.  When  a  mixture  of 
several  gases  is  maintained  over  a  liquid,  each  dissolves  independently 
of  the  other  and  in  accordance  with  its  own  partial  pressure.  This 
is  known  as  the  law  of  Dalton,  and  its  statement  was  the  outcome  of 
some  of  his  earlier  experiments,  made  in  1803.  For  example,  when 
air  is  inclosed  over  water  at  760  mm.  pressure,  approximately  ^  of 
this  pressure,  namely,  152  mm.,  is  due  to  oxygen,  and  |,  or  608  mm., 
is  due  to  nitrogen  and  other  gases.  According  to  Dalton's  law  the 
quantity  of  oxygen  dissolving  in  water  exposed  to  air  will  be  the  same 
as  if  the  water  were  confined  in  a  space  containing  only  oxygen  at  a 


126  GENERAL  CHEMISTRY 

pressure  of  152  mm.  When  a  solution  of  a  gas  other  than  nitrogen 
or  oxygen  is  exposed  to  the  open  air,  the  gas  will  in  general  escape, 
for  there  is  no  opposing  partial  pressure  to  keep  it  in  solution. 

It  follows  that  when  one  gas  is  bubbled  for  some  time  through  a 
solution  of  a  second,  the  latter  will  be  washed  out  of  the  solution,  for 
in  the  atmosphere  over  the  liquid  it  exerts  no  opposing  pressure. 
When  a  solution  of  a  gas  is  boiled,  water  vapor  takes  the  place  of 
this  displacing  gas.  It  bubbles  through  the  solution,  maintaining  an 
atmosphere  of  steam  over  the  liquid,  and  so  prevents  a  pressure  by  the 
dissolved  gas.  Hence,  gases  are  in  general  insoluble  in  boiling  solvents. 

If,  however,  the  gas  is  a  very  soluble  one,  and  it  happens,  in  the 
process  of  boiling,  that  the  ratio  by  weight  of  the  gas  to  the  steam  in 
the  atmosphere  over  the  solution  comes  to  be  the  same  as  the  ratio  of 
the  dissolved  gas  to  the  water  in  the  solution,  then  the  partial  pres- 
sure of  the  gas  will  be  steadily  maintained,  and  the  solution  will 
boil  without  any  change  in  concentration.  This  happens  in  a  number 
of  familiar  instances.  For  example,  a  solution  of  hydrogen  chloride 
boils  with  constant  concentration  at  110°,  when  it  contains  20.24  per 
cent  of  the  gas.  The  gas  cannot  be  boiled  out  of  such  a  solution. 

Solutions  of  liquids  in  liquids.  Two  liquids  may  conduct  themselves 
toward  each  other  in  either  of  two  ways :  they  may  be  freely  soluble 
(or  miscible)  in  all  proportions  or  each  may  reach  a  definite  limit  of 
saturation  with  the  other. 

Freely  miscible  liquids.  A  number  of  familiar  liquids  mix  freely 
with  water  in  all  proportions,  among  them  being  alcohol,  glycerin, 
and  many  acids,  such  as  nitric,  sulf uric,  and  acetic.  Most  oils  and  fats 
are  also  miscible  with  each  other  in  all  proportions.  In  some  cases,  as 
with  alcohol  and  water,  each  liquid  appears  to  retain  its  own  chemical 
characteristics  in  solution ;  in  other  cases,  especially  with  acids  dis- 
solved in  water,  new  chemical  characteristics  are  acquired  along  with 
the  old  ones,  suggesting  the  existence  of  some  such  chemical  equilib- 
rium as  that  described  in  connection  with  ammonia.  Several  properties 
of  liquid  solutions  are  of  importance. 

1.  Vapor  pressure.  Before  mixing,  each  liquid  has  its  own  charac- 
teristic vapor  pressure  at  the  temperature  in  question.  After  mixing, 
it  is  found  that  each  liquid  has  diminished  the  vapor  pressure  of  the 
other,  so  that  the  vapor  pressure  of  the  solution  is  never  as  great  as 
the  sum  of  the  two  original  ones.  It  may  be  greater  or  less  than  that 
of  either  liquid  taken  separately,  or  it  may  have  an  intermediate  value. 


SOLUTIONS  127 

It  also  depends  on  the  relative  concentrations  of  the  two  liquids.  In 
any  case  the  vapor  from  the  solution  at  any  definite  temperature  will 
be  a  mixture  of  that  of  each  liquid,  provided  each  has  a  sensible  vapor 
pressure  at  that  temperature. 

2.  Boiling  point.  On  heating  a  solution  of  one  liquid  in  another, 
the  total  vapor  pressure  increases,  and  when  it  just  exceeds  the  oppos- 
ing pressure  of  the  atmosphere,  the  solution  boils.    From  what  has 
been  said  in  regard  to  the  vapor  pressure  of  solutions  it  will  be  seen 
that  we  can  form  little  idea  as  to  the  boiling  point  of  a  solution  from 
the  known  boiling  points  of  its  constituents.    It  may  be  lower  or  higher 
than  that  of  either  constituent.    Usually  it  has  an  intermediate  value. 
Since  the  relative  weights  of  the  two  vapors  escaping  from  the  solu- 
tion are  not  in  general  the  same  as  those  of  the  two  liquids  constituting 
the  solution,  the  composition  of  the  liquid  will  vary  during  the  proc- 
ess of  boiling.    This  will  lead  to  variations  in  the  composition  of  the 
vapor,  and  a  steady  change  in  boiling  point.    A  solution,  therefore, 
has  in  general  no  constant  boiling  point. 

3.  Fractional  distillation.  The  liquid  having  the  greater  vapor  pres- 
sure will  in  general  pass  away  from  the  solution  more  rapidly  than 
the  one  of  lower  vapor  pressure  (higher  boiling  point).    If  the  vapors 
are  condensed,  as  described  in  the  distillation  of  water,  and  the  re- 
sulting liquid  is  collected  in  consecutive  portions  by  changing  the 
receiver  at  intervals,  the  first  portions  will  be  richer  in  the  more  vola- 
tile constituent,  the  higher-boiling  liquid  being  largely  obtained  in 
the  latter  portions.    By  repeating  the  process  with  each  portion  ob- 
tained in  the  first  operation,  the  two  liquids  may  in  time  be  separated 
from  each  other.    Such  a  process  is  called  fractional  distillation.    It  is 
of  the  greatest  importance  in  many  industries,  such  as  the  refining  of 
petroleum  and  the  manufacture  of  alcohol  and  glycerin.    We  have 
already  seen  that  relatively  pure  oxygen  is  prepared  from  liquid  air 
in  this  way. 

4.  Constant-boiling  solutions.   It  occasionally  happens  that  there  is 
one  particular  concentration  of  a  solution  which  has  a  lower  vapor 
pressure  (higher  boiling  point)  than  any  other  concentration,  or  than 
that  of  either  constituent  taken  separately.    When  such  a  solution  is 
distilled,  one  or  the  other  constituent  vaporizes  first,  the  concentra- 
tion tending  toward  that  of  lowest  vapor  pressure.    When  this  concen- 
tration is  reached,  the  solution  boils  with  constant  boiling  point  like 
a  pure  substance,  and  the  distillate  has  the  same  composition  as  the 


128  GENERAL  CHEMISTRY 

solution  remaining  in  the  still.  Such  a  constant-boiling  solution  can- 
not be  altered  in  composition  by  repeated  distillation.  An  example 
of  such  conduct  is  found  in  aqueous  solutions  of  nitric  acid,  the  con- 
stant-boiling solution  consisting  of  68  per  cent  acid  and  32  per  cent 
water.  The  constant  boiling  solution  of  sulfuric  acid  contains  98.33 
per  cent  acid. 

Sparingly  miscible  liquids.  If  water  is  shaken  for  a  few  moments  with  ether 
or  chloroform  in  a  suitable  vessel,  such  as  a  separatory  funnel  (Fig.  34,  p.  70), 
and  set  aside  for  a  short  time,  the  more  or  less  turbid  liquid  gradually  clears 
and  two  liquid  layers  form,  the  one  of  smaller  density  floating  on  the  heavier. 
Each  of  these  layers  is  a  saturated  solution.  In  the  case  of  ether  and  water  the 
upper  layer  consists  of  ether  saturated  with  water ;  the  lower,  of  water  saturated 
with  ether.  With  some  pairs  of  liquids,  especially  if  the  densities  are  not  greatly 
different,  this  parting  into  two  layers  is  very  slow.  The  liquid  remains  turbid 
and  is  called  an  emulsion,  the  name  suggesting  a  conspicuous  example,  namely, 
milk.  Ultimately  the  separation  takes  place,  the  lighter  solution  rising  to  the  top. 

Solutions  of  solids  in  liquids.  A  solid  dissolved  in  a  liquid  is  by  far 
the  most  familiar  type  of  solution.  In  the  following  discussion  it 
should  be  remembered  that  we  are  dealing  with  true  solutions  only. 
Thus  it  is  sometimes  said  that  zinc  dissolves .  in  hydrochloric  acid. 
In  this  case,  however,  the  solution  is  preceded  by  an  undoubted  chem- 
ical reaction  whereby  the  zinc  is  converted  into  zinc  chloride,  and  it 
is  this  compound  which  is  obtained  when  the  solution  is  evaporated  to 
dryness.  With  solutions  such  as  we  are  now  considering,  evaporation 
leaves  the  solute  in  its  original  chemical  condition. 

Molar  solutions.  In  stating  the  concentration  of  a  solution  we  may 
obviously  make  use  of  the  percentage  system.  It  is  often  more  desir- 
able to  state  the  number  of  formula  or  molecular  weights  (measured 
in  grams)  which  a  given  volume  of  the  solution  contains.  When  as 
many  grams  of  a  substance  as  there  are  units  in  its  molecular  weight 
is  dissolved  so  as  to  make  a  liter  of  solution,  it  is  said  to  be  a  molar 
or  gram-molecular  solution.  Thus,  a  molar  solution  of  sodium  hy- 
droxide (NaOH)  contains  40.01  g.  of  the  compound  in  1  1.,  while 
that  of  nitric  acid  (HNO3)  contains  63.02  g.  If  15  g.  of  sodium 
hydroxide  is  dissolved  so  as  to  form  1  1.  of  solution,  the  molar  con- 
centration is  0.375  (15  -*-  40.01). 

Saturated  solutions.  When  a  lump  of  sugar  is  placed  in  a  small 
beaker  and  covered  with  water,  as  represented  in  Fig.  58,  it  gradually 
diminishes  in  size  and  passes  into  solution,  particles  leaving  it  and  dif- 
fusing through  the  solvent.  If  there  is  enough  sugar,  and  a  long  enough 


SOLUTIONS  129 

time  elapses,  the  concentration  of  the  sugar  in  the  solution  reaches  a 
definite  limiting  value,  and  we  say  that  the  sugar  ceases  to  dissolve  and 
that  the  solution  is  saturated.  There  is  good  reason  for  thinking  that 
particles  continue  to  leave  the  lump,  but  that  an  equilibrium  has  been 
reached,  the  rate  of  departure  of  the  particles  being  equal  to  the  rate 
of  their  return.  A  saturated  solution  may  there- 
fore be  defined  as  one  which  is  in  equilibrium 
with  the  undissolved  solute. 


It  takes  a  very  long  time  for  a  solution  to  become 
saturated  in  this  way.  If  a  few  crystals  of  the  highly 
colored  salt  potassium  permanganate  are  placed  in  the 
bottom  of  a  tall  cylinder,  which  is  then  filled  with  water  FIG.  58 

in  such  a  way  as  not  to  disturb  them,  and  the  cylinder 

set  aside  in  a  quiet  place,  months  will  pass  before  the  color  of  the  solution 
becomes  uniform.  If  the  solvent  is  stirred  or  shaken  vigorously  during  solution, 
saturation  may  be  reached  in  a  short  time.  The  rate  at  which  substances  dis- 
solve and  reach  saturation  varies  greatly,  and  it  does  not  bear  any  very  simple  rela-  • 
tion  to  the  quantity  ultimately  going  into  solution.  It  is  dependent  on  the  specific 
properties  of  both  solute  and  solvent,  on  the  rate  of  diffusion  of  the  dissolved 
substance,  on  its  fineness  of  division,  and  doubtless  on  many  other  conditions. 

That  the  concentration  of  a  saturated  solution  at  a  given  tempera- 
ture has  a  definite  value  may  be  shown  by  obtaining  saturation  in  still 
a  different  way.  Most  solids  are  more  soluble  in  hot  liquids  than  in 
cold.  If  we  approximately  saturate  a  solution  at  a  higher  temperature 
and  then  reduce  the  temperature  to  the  desired  point,  taking  care  to 
have  some  of  the  solid  present  all  the  tune,  the  excess  of  the  solute 
crystallizes  out  and  almost  at  once  the  solution  comes  to  the  same 
concentration  as  was  reached  in  the  other  slower  way. 

Supersaturated  solutions.  If  we  neglect  to  have  any  of  the  solid 
present  during  the  cooling,  it  may  happen  that  the  excess  of  solute 
will  fail  to  crystallize  out.  The  solution  will  then  contain  more  than 
the  normal  saturation  quantity  of  the  solute,  and  is  said  to  be  super- 
saturated. This  is  apt  to  occur  in  the  case  of  very  soluble  solids,  with 
salts  containing  much  water  of  crystallization,  —  for  example,  with  so- 
dium sulfate  (Na^SCVlO  H2O),  sodium  thiosulfate  (Na^Og  •  5  H2O), 
and  ferric  nitrate  (Fe(NO3)3  •  9  H2O),  —  and,  in  general,  in  all  cases 
where  the  solution  becomes  sirupy.  It  must  be  remembered  that 
saturation  is  normally  an  equilibrium.  When  one  member  of  the 
equilibrium  is  absent,  abnormal  results  may  be  expected.  The  intro- 
duction of  even  the  smallest  fragment  of  the  solid  solute  into  a 


130 


GENEBAL  CHEMISTRY 


supersaturated  solution  will  bring  about  the  crystallization  of  the 
excess  of  solute.  When  this  is  thrown  out  of  solution  very  rapidly, 
it  appears  in  a  very  finely  divided  form  called  a  precipitate. 

Determination  of  solubility.  The  preceding  discussion  suggests  a  ready 
method  for  determining  the  solubility  of  a  solid.  A  saturated  solution  is  pre- 
pared, either  by  long-continued  stirring  of  the  solid  with  the  solvent  at  the 
desired  temperature  or  by  approximately  saturating  the  solvent  at  a  higher  tem- 
perature and  cooling  in  the  presence  of  the  solid.  A  measured  portion  of  the 
solution  is  then  filtered  off,  evaporated  to  dryness,  and  the  residue  weighed. 

Effect  of  temperature  on  solubility.  Change  of  temperature  always 
changes  the  solubility  of  a  solute.  As  a  rule,  solids  are  more  soluble 
in  hot  than  in  cold  solvents,  though  occasionally  the  reverse  is  true. 
Many  compounds  of  calcium,  including  the  hydroxide  Ca(OH)2,  are 
of  this  latter  class,  while  the  solubility  of  common  salt  is  very  little 
affected  by  the  temperature.  The  table  shows  the  change  in  solubility 
•in  a  few  familiar  cases. 

TABLE  OF  SOLUBILITY  OF  SOLIDS 


SUBSTANCE 

FORMULA 

WEIGHT  DISSOLVED  BY  100  cc.  OF  WATER  AT 

0° 

20° 

100° 

Calcium  chloride     . 

CaCl2 

59.5  g. 

74.5  g. 

159.0  g. 

Sodium  chloride 

NaCl 

35.70  g. 

36.0  g. 

39.80  g. 

Potassium  nitrate    . 

KN03 

13.30  g. 

31.6  g. 

246.0  g. 

Copper  sulfate     . 

CuSO4 

14.30  g. 

21.7  g. 

75.4  g. 

Calcium  sulfate  . 

CaSO4 

0.759  g. 

0.203  g. 

0.162  g. 

Calcium  hydroxide  . 

Ca(OH)2 

0.185  g. 

0.165  g. 

0.077  g. 

Solubility  curves.  As  a  rule,  the  solubility  of  a  solute  does  not  vary 
with  temperature  in  any  regular  manner,  so  we  can  best  represent  the 

facts  in  a  given  case  by 
a  curve,  plotting  the 
temperature  as  the  ab- 
scissa and  the  concen- 
tration as  the  ordinate. 
The  diagram  (Fig.  59) 
shows  a  few  typical 
curves,  the  concentra- 

40"  so9  60°   70°  60°  90"  too'T    tion  being  measured  by 
FIG.  59  the  number  of  grams 


too 


50 


10° 


30" 


SOLUTIONS 


131 


of  solute  dissolved  in  100  g.  of  solvent.  It  will  be  seen  that  some 
substances  are  very  soluble,  while  others  are  not.  The  solubility  of 
some  increases  rapidly  with  temperature,  while  in  other  cases  the 
increase  is  small. 

Breaks  in  solubility  curves.  The  curve  for  sodium  sulfate  (Na2SO4-10  H2O) 
presents  a  singular  feature  (Fig.  60).  Up  to  32.4°  it  rises  rapidly  in  a  normal  way, 
but  above  that  point  it  steadily  falls.  An  examination  of  the  solids  separating 
above  and  below  this  point  shows  that  they  are  not  the  same.  Below  32.4°  the  solid 
is  Na2SO4  •  10  H2O,  while  above  this  point  it  is  the  anhydrous  salt  Xa2SO4.  The 
point  32.4°,  therefore,  marks  the  temperature  at  which  the  one  salt  changes  into 
the  other ;  like  the  freezing  point  of  a  liquid,  it  is  called  a  transition  point. 
Sharp  breaks  like  this  in  a  solubility  curve  always  suggest  some  chemical  trans- 
formation in  the  solute, 
so  that  the  study  of  such 
curves  is  of  much  impor- 
tance in  giving  informa- 
tion as  to  the  changes 
taking  place  in  solution. 

Solutions  of  solids  in 
solids.  A  number  of 
cases  have  been  noted  in 
which  one  solid  slowly 
diffuses  into  another, 
much  as  sugar  diffuses 
into  water.  Gold  will 
diffuse  into  lead  when 
clean  pieces  of  the  two  metals  are  tightly  clamped  together ;  copper- 
plated  zinc  gradually  turns  lighter  in  color  from  mutual  diffusion  of 
the  two  metals.  The  products  in  such  cases  are  called  solid  solutions. 
Similar  solutions  can  be  prepared  in  special  cases  by  melting  the  two 
solids  together.  If  they  form  a  liquid  solution,  they  may  freeze  as  a 
homogeneous  solid  with  many  of  the  properties  of  a  true  solution.  In 
other  cases  they  do  not  do  this ;  one  solid  or  the  other  separates  as  a 
pure  material,  or  definite  compounds  form. 

Solutions  of  amorphous  bodies.  It  will  be  recalled  that  in  Chapter  VI 
amorphous  bodies,  such  as  glass,  waxes,  and  glue,  were  classed  as 
liquids  rather  than  as  solids.  It  is  of  interest  to  note  that  such  bodies 
do  not,  as  a  rule,  have  a  definite  solubility.  They  are  apt  to  mix  with  a 
solvent  in  all  proportions,  as  is  true  of  glue,  the  solution  gradually 
drying  up  to  a  semisolid  form  as  the  solvent  is  evaporated.  This  con- 
duct is  more  like  that  of  one  liquid  dissolved  in  another  than  of  a  solid 


20 


10 


132  GENERAL  CHEMISTEV 

dissolved  in  a  liquid.  Bodies  forming  solutions  of  this  kind  are  called 
colloids,  from  the  Greek  word  meaning  "glue,"  and  their  solutions  are 
called  colloidal  solutions.  Glues,  pastes,  and  jellies  belong  to  this  class. 
Colloids.  Colloidal  solutions  of  many  substances  normally  crystalline 
can  be  prepared  in  special  ways.  Thus,  when  the  points  of  two  plati- 
num wires  are  dipped  under  water,  as  shown  in  Fig.  61,  and  an  electric 
arc  is  formed  between  them  (by  a  current  of  about  30  volts  and  10 
amperes),  particles  of  platinum  are  torn  off  from  the  negative  wire, 

and  the  water  turns  to  a  brownish  color. 
Other  metals,  such  as  gold,  silver,  and 
copper,  act  in  a  similar  way,  giving  highly 
colored  solutions.  Solutions  of  the  same 
character  can  be  prepared  by  treating  sol- 
uble compounds  of  these  metals,  under 
special  conditions,  with  reagents  which 
normally  precipitate  the  metal.  Indeed, 
a  great  number  of  reactions  which  usually 

yield  precipitates  will,  under  special  conditions,  yield  a  colloidal  solu- 
tion, so  that  it  appears  that  the  formation  of  an  amorphous  colloid  is 
often  the  first  step  in  precipitation,  the  amorphous  body  gradually 
becoming  crystalline  and  appearing  as  a  true  precipitate.  When  the 
substance  is  capable  of  crystallization,  the  colloid  is  not  a  stable  form, 
and  in  such  cases  it  is  never  permanent.  In  many  cases  special  methods 
have  shown  that  with  such  bodies  the  particles  are  merely  in  a  very 
finely  divided  condition  in  suspension  rather  than  in  solution,  and  the 
term  colloidal  suspension  is  often  employed.  There  is  apparently  a  con- 
tinuous gradation  from  ordinary  suspensions,  such  as  chalk  in  water, 
to  colloidal  suspensions  which  never  settle,  thence  to  colloidal  solu- 
tions in  which  no  suspended  particles  can  be  detected,  and  finally  to 
ordinary  solutions  of  crystallized  compounds,  and  no  sharp  lines  can 
be  drawn  between  these  classes. 

Regularities  observed  in  solutions.  A  careful  study  of  the  physical 
properties  of  solutions,  particularly  those  of  solids  in  liquids,  has 
brought  to  light  regularities  which  have  led  to  the  development  of 
certain  theories  regarding  the  nature  of  solutions.  These  regularities 
are  confined  to  dilute  solutions,  and  they  do  not  hold  good  when  the 
solute  belongs  to  either  of  the  three  classes  of  compounds  known  as 
acids,  bases,  and  salts,  or,  collectively,  as  electrolytes.  A  discussion  of 
solutions  of  electrolytes  will  be  postponed  until  the  following  chapter- 


SOLUTIONS 


133 


1.  Freezing  point  of  solutions.  It  has  been  known  for  a  very  long 
time  that  the  freezing  point  of  a  solution  is  lower  than  that  of  the 
pure  solvent.  Doubtless  one  of  the  first  observations  of  this  kind 
was  that  salt  water  remains  unfrozen,  while  fresh  water  near  by 
freezes.  Since  the  time  of  Blagden,  who  was  secretary  to  Caven- 
dish, it  has  been  known  that  the  lowering  of  the  freezing  point  of 
a  solution  is  proportional  to  the  concentration.  For  example,  a 
2  per  cent  solution  of  sugar  freezes  twice  as  much  below  zero  as 
does  a  1  per  cent  solution,  and  a  3  per  cent  solution  three  times  as 
much  below  zero. 

If,  instead  of  measuring  the  concentrations  in  percentages,  we  express 
them  in  molar  concentrations,  most  interesting  results  are  obtained. 
It  is  found  that  the  lowering  of  the  freezing  point  of  a  definite  weight 
of  a  solvent,  such  as  water,  produced  by  a  molar  weight  of  glycerin 
and  of  sugar  is  the  same.  Equimolecular  weights  of  all  substances 
(aside  from  electrolytes),  irrespective  of  their  chemical  properties,  pro- 
duce an  equal  lowering  of  the  freezing  point  of  a  definite  weight  of  a 
given  solvent,  within  the  limits  of  unavoidable  error.  This  generaliza- 
tion is  known  as  the  law  of  Raoult  and  was  formulated  in  1883. 
The  following  table  illustrates  this  law : 

LOWERING  OF  THE   FREEZING  POINT  OF  WATER 


SOLUTE 

FORMULA 

LOWERING  PRODUCED 
ix  1000  cc.  OF  WATER 
BY  1  GRAM-MOLECULE 
OF  SUBSTAXCE 

jMcthvl  alcohol 

CH4O 
C2H60 

C12H22O11 

^If^MPll 

C3H60 

1.90° 
1.87° 
2.02° 
2.06° 
1.92° 

Ethyl  alcohol 

Can6  suffar   

Milk  sugar    .     .     .     .  f.     .     . 
A.C6ton6 

2.  Lowering  of  the  vapor  pressure.  Raoult  found  that  a  similar  law 
holds  for  the  lowering  of  the  vapor  pressure  of  a  solvent.  If  molar 
weights  of  solutes  which  are  neither  electrolytes  nor  liquids  with 
sensible  vapor  pressures  are  dissolved  in  a  definite  weight  of  a  solvent, 
the  vapor  pressure  of  the  solvent  is  lowered  equally  in  all  cases.  The 
lowering  is  independent  of  the  chemical  properties  of  the  substances, 
but  depends  merely  upon  the  numerical  ratio  of  the  two  kinds  of 
molecules  which  constitute  the  solution. 


134  GENERAL  CHEMISTEY 

3.  Elevation  of  the  boiling  point.    Since  the  boiling  point  of  a  liquid 
is  the  temperature  at  which  its  vapor  pressure  just  exceeds  the  oppos- 
ing pressure  of  the  atmosphere,  it  is  evident  that  any  condition  which 
lowers  the  vapor  pressure  will  raise  the  boiling  point,  since  the  liquid 
will  have  to  be  heated  higher  to  return  to  the  original  vapor  pressure. 
Lowering  of  the  vapor  pressure  is  found  to  be  almost  exactly  pro- 
portional to  rise  of  boiling  point.   It  follows  that  molar  weights  of  any 
substance  not  an  electrolyte  and  not  having  a  vapor  pressure  of  its 
own,  when  dissolved  in  a  definite  weight  of  a  solvent,  occasion  the 
same  rise  in  the  boiling  point  of  the  solvent, 

4.  Osmotic  pressure.  We  have  seen  that  a  substance  in  solution  tends 
to  distribute  itself  uniformly  throughout  the  entire  volume  of  the 
solvent.    The  force  which  brings  about  this  distribution  is  called  os- 
motic pressure.    We  do  not,  as  a  rule,  notice  the  existence  of  any  such 
force  or  gain  any  idea  as  to  its  magnitude,  because  it  is  only  when  dif- 
fusion can  be  prevented  that  the  force  becomes  apparent.    This  can 
be  accomplished  by  separating  a  solution  from  some  of  the  pure  sol- 
vent by  means  of  a  membrane  through  which  the  solvent  can  readily 
pass,  but  which  restrains  the  dissolved  molecules.    Such  a  membrane 
is  said  to  be  semipermeable.    Many  animal  and  vegetable  membranes 
are  of  this  sort,  so  that  osmotic  pressure  plays  an  important  part 
in  physiological  processes. 

Qualitative  demonstration  of  osmotic  pressure.  A  piece  of  parchment  paper  is 
fastened  tightly  over  the  bell  of  a  funnel  tube  by  means  of  windings  of  cord  and 
sealing  wax  (Fig.  62).  A  moderately  concentrated 
solution  of  sugar  is  placed  in  the  vessel  so  constructed, 
which  is  then  immersed  in  pure  water.  The  sugar 
cannot  pass  through  the  parchment  to  distribute  itself 
throughout  the  water,  but  the  water  passes  into  the 
bell,  the  solution  rising  in  the  tube  until  the  force 
of  osmotic  pressure  is  balanced  by  the  hydrostatic 
pressure  of  the  column  of  liquid. 

Quantitative  measurements.  An  apparatus  such  as 
the  one  just  described  is  not  adapted  to  the  exact 
measurement  of  osmotic  pressure,  because  the  mem- 
brane is  not  strong  enough  to  withstand  the  pressure 
which  develops.  It  has  been  found  that  a  very  satis- 
factory substitute  is  obtained  by  precipitating  the 
IG*  compound  known  as  copper  ferrocyanide  in' the  pores 

of  a  fine-grained,  unglazed  porcelain  cell  similar  to 

a  small  battery  jar.    This  compound  is  an  amorphous  body  which  is  semipermea- 
ble, and  the  walls  of  the  minute  pores  within  which  it  is  precipitated  give  it  the 


SOLUTIONS 


135 


requisite  strength.  Morse  has  prepared  very  perfect  cells  of  this  kind  by  precip- 
itating the  copper  ferrocyanide  by  electrolytic  methods,  the  cells  withstanding 
a  pressure  of  as  much  as  100  atmospheres. 

Such  a  cell  is  filled  with  a  solution  whose  osmotic  pressure  is  to  be  measured ; 
the  cell  is  tightly  connected  with  a  pressure  gauge,  or  manometer,  and  is  then  im- 
mersed in  pure  water.    The  water  tends  to  flow  into  the 
cell  with  a  force  equal  to  the  osmotic  pressure,  and  this 
is  registered  on  the  gauge.    Fig.  63  shows  the  apparatus 
recently  employed  by  Morse,  A  being  the  cell  and  B  the 
manometer. 


B 


The  laws  of  osmotic  pressure.  The  first  quan- 
titative measurements  of  osmotic  pressure  were 
made  by  the  German  botanist  Pfeffer  in  1877, 
but  the  recent  work  of  Morse  in  America  and 
the  Earl  of  Berkeley  in  England  are  of  a  very 
much  higher  order  of  accuracy.  Their  results 
show  that  within  the  limits  of  unavoidable  error 
the  osmotic  pressure  is  proportional  to  the  molec- 
ular concentration  of  the  solution,  and  does  not 
depend  upon  the  character  of  the  solute,  provided 
it  is  not  an  electrolyte.  It  is  also  proportional  to  the  absolute 
temperature,  so  that  the  osmotic  pressure  exerted  by  a  substance  in 
solution  is  equal  to  the  gas  pressure  which  it  would  exert  if  it 
were  a  gas  occupying  the  same  volume  as  that  of  the  solvent.  The 
table  gives  some  of  Morse's  results,  the  last  column  showing  how 
closely  the  two  pressures  agree. 


FIG.  63 


OSMOTIC  PRESSURE  OF  CANE  SUGAR 


MOLECULAR  CON- 
CENTRATION 

AT  20° 

OSMOTIC  PRESSURE 
(Atmospheres) 

CALCULATED  GAS 
PRESSURE 
(Atmospheres) 

RATIO  OF  OSMOTIC 
PRESSURE  TO  GAS 
PRESSURE 

0.1 

2.590 

2.390 

1.084 

0.2 

5.064 

4.78 

1.062 

0.3 

7.605 

7.17 

1.060 

0.4 

10.137 

9.56T 

1.060 

0.5 

12.748 

11.95 

1.067 

0.6 

15.388 

14.339 

1.073 

0.7 

18.127 

10.729 

1.084 

0.8 

20.905 

19.119 

1.093 

0.9 

23.717 

21.509 

1.103 

1.0 

26.638 

23.899 

1.115 

136  GENERAL  CHEMISTKY 

Summary.  From  the  foregoing  paragraphs  it  appears  that  three  of 
the  most  important  of  the  physical  constants  of  a  solvent,  namely, 
the  freezing  and  boiling  points  and  the  vapor  pressure,  are  changed 
to  an  equal  degree  by  molar  quantities  of  a  dissolved  substance,  and 
that  it  is  the  relative  number  of  the  two  kinds  of  molecules,  not  their 
character,  which  determines  the  extent  of  the  change.  Furthermore, 
the  osmotic  pressure  of  a  solution  is  governed  by  the  same  numerical 
ratio  and  is  independent  of  the  character  of  the  substances. 

The  picture  which  this  presents  to  us  is  quite  similar  to  our  kinetic 
idea  of  gases.  The  various  kinds  of  molecules  appear  to  distribute 
themselves  throughout  the  solvent  much  as  a  gas  distributes  itself  in 
the  space  in  which  it  is  inclosed.  The  osmotic  pressure  which  they 
produce  is  quite  analogous  to  the  pressure  of  a  gas,  and  the  other 
effects  are  dependent  merely  upon  the  number  of  molecules  present 
in  a  given  volume  of  the  solvent,  just  as  the  effects  produced  by 
a  gas  are  dependent  upon  the  number  confined  in  a  given  space. 


CHAPTER  XII 

IONIZATION  AND  ELECTROLYSIS 

Introduction.  In  the  preceding  chapter  it  has  been  shown  that  the 
effect  produced  upon  such  properties  of  a  solution  as  the  boiling  and 
freezing  points,  the  vapor  pressure,  and  the  osmotic  pressure  depends 
upon  the  molecular  ratio  of  the  two  components  of  the  solution  (the 
solvent  and  the  solute)  and  not  upon  their  specific  chemical  properties, 
just  as  gas  pressure  depends  upon  the  number  and  not  the  kind  of 
molecules  present.  The  fact  has  been  emphasized  that  acids,  basesr 
and  salts  do  not  exhibit  these  regularities  but  give  abnormal  results. 
A  few  statements  regarding  the  nature  of  these  classes  of  compounds 
will  be  of  advantage  before  proceeding  to  a  discussion  of  their  excep- 
tional character. 

The  acids  include  a  great  variety  of  compounds,  all  containing 
hydrogen,  which,  under  suitable  conditions,  can  be  replaced  by  various 
metals.  They  also  change  the  color  of  certain  substances  called  indi- 
cators ;  thus,  the  indicator  known  as  blue  litmus  turns  red  in  the 
presence  of  an  acid.  The  most  common  acids  are  hydrochloric  (HC1), 
nitric  (HNOg),  and  sulfuric  (H^OJ. 

The  inorganic  bases  are  compounds  which  are  composed  of  some 
metallic  element  combined  with  oxygen  and  hydrogen,  and  include 
such  substances  as  potassium  hydroxide  (KOH)  and  calcium  hydroxide 
(Ca(OH)2).  The  soluble  bases  also  produce  color  changes  in  indi- 
cators, turning  red  litmus  blue. 

Salts  are  formed  by  the  action  of  a  base  upon  an  acid.  The  hydro- 
gen of  the  acid  combines  with  the  oxygen  and  hydrogen  of  the  base, 
forming  water,  and  the  other  constituents  unite  to  form  a  salt.  The 
following  equations  will  illustrate  this  reaction : 

KOH  +  HC1  =  HO  +  KC1 

(base)  (acid)  (salt) 

Ca(OH)2  +  H  S04  =  2  H  O  +  CaSO4 

(base)  (acid)  (salt) 


Irregularities  in  solutions  of  electrolytes.  When  an  acid,  a  base,  or  a 
salt  is  dissolved  in  water,  the  effect  produced  upon  the  freezing  or 

137 


138 


GENERAL  CHEMISTRY 


boiling  point  or  the  vapor  pressure  of  the  water,  and  also  the  osmotic 
pressure  of  the  solution,  is  greater  than  that  calculated  from  the  laws 
developed  in  the  last  chapter.  The  deviation  from  the  calculated 
effect  varies  between  wide  limits.  With  some  electrolytes  the  excess 
is  slight ;  with  others  it  reaches  a  double  value.  Moreover,  it  varies 
with  the  concentration,  being  relatively  greater  in  dilute  solutions  than 
in  more  concentrated  ones.  The  table  shows  the  effects  produced  upon 
the  freezing  point  of  water  by  a  typical  salt,  base,  and  acid  when 
compared  with  that  produced  by  alcohol,  which  is  not  an  electrolyte. 

LOWERING   OF   THE   FREEZING   POINT   OF*WATER 


SUBSTJLNCE 

CHARACTER 

FORMULA 

LOWERIXG  PRODUCED  IX  1000  CC. 
OF  WATER  BY  1  GRAM-MOLE- 
CULE OF  SUBSTANCE 

Alcohol               .... 

nonelectrol  vte 

C9H.O 

1.872 

Hydrogen  chloride      .     . 

acid 

HC1 

3.806 

Hydrogen  nitrate    .     .     . 

acid 

HNO3 

4.100 

Potassium  hydroxide  .     . 

base 

KOH 

3.773 

Sodium  hydroxide  . 

base 

NaOH 

3.465 

Sodium  chloride     .     .     . 

salt 

NaCl 

3.546 

Potassium  nitrate  . 

salt 

KX03 

2.729 

If  the  conclusion  reached  in  the  study  of  nonelectrolytes  is  correct, 
—  namely,  that  it  is  the  relative  number  of  molecules  of  the  solute,  not 
its  character,  which  determines  the  effect,  —  then  in  the  case  of  elec- 
trolytes the  thought  occurs,  May.  it  not  be  that,  when  dissolved  in 
water,  electrolytes  undergo  a  dissociation  into  units  smaller  than  mole- 
cules, which  therefore  increases  the  number  of  independent  particles  ? 
We  are  familiar  with  such  a  dissociation  occasioned  by  heat.  It  will 
be  remembered  that  at  high  temperatures  water  dissociates  to  some 
extent,  setting  up  the  equilibrium  represented  in  the  equation 


the  number  of  independent  units  increasing  from  2  to  3  in  so  far  as 
the  reaction  proceeds,  and  the  gas  pressure  increasing  correspondingly. 
We  shall  first  seek  for  evidence  that  there  is  a  somewhat  similar  dis- 
sociation in  solutions  of  electrolytes,  and  then  endeavor  to  form  a 
theory  as  to  its  character. 

Evidence  for  dissociation  in  solutions  of  electrolytes.  There  are  a 
number  of  lines  of  evidence  which  may  be  mentioned  in  support  of 
the  view  that  in  solutions  of  electrolytes  a  dissociation  takes  place. 


N^  AND  ELECTROLYSIS  139 

1.  Abnormal  constants  of  solutions  of  electrolytes.  The  fact  of  the  ex- 
cessive effect  produced  by  electrolytes  on  the  lowering  of  the  freezing 
point  and  vapor  pressure,  on  the  raising  of  the  boiling  point,  and  on 
the  magnitude  of  osmotic  pressure  is  evidence  pointing  toward  the 
increase  in  number  of  independent  units  in  solution,  provided  our 
conclusions  in  regard  to  nonelectrolytes  are  justified. 

2.  Chemical  reactions  of  solutions  of  electrolytes.  There  is  a  simplicity 
in  the  chemical  action  of  electrolytes  upon  each  other  in  solution 
which  is  strongly  suggestive  of  the  presence  of  dissociated  products. 
The  prevailing  type  of  reaction  in  their  solutions  is  illustrated  by 
such  equations  as  the  following: 


Ba(X  +  H2SO4  =  BaSO4  4-  H«,O2 

KOH  +  HNOS  =  KX03  +  HOH 

AgNO3  4-  NaCl  =  AgCl  +  NaXO, 


It  will  be  seen  that  these  reactions  express  a  simple  interchange,  the 
metals  exchanging  places  with  each  other  or  with  hydrogen,  while 
such  groups  as  SO4  and  NO8  go  through  the  reaction  as  units.  More- 
over, the  reactions  proceed  with  great  rapidity,  indicating  that  little 
preliminary  work  is  involved  in  decomposing  the  materials  into  these 
groups.  This  is  very  different  from  most  of  the  reactions  of  nonelec- 
trolytes in  solution,  which  are  slow,  or  from  those  occasioned  by  heat, 
such  as  the  decomposition  of  potassium  chlorate,  the  course  of  which 
in  general  cannot  be  predicted  at  all. 

It  would  appear  that  with  electrolytes  there  is  a  distinct  line  of 
weakness  in  the  molecule  —  that  sulfuric  acid  tends  to  .part  into  hydro- 
gen and  the  group  SO4  with  great  ease,  while  the  group  SO4  tends  to 
act  as  a  unit.  This  suggests  that  in  solution  the  acid  may  really  be  in 
the  condition  represented  by  the  equilibrium  equation 

H2S04:<=±2H  +  S04 

3.  The  law  of  thermoneutrality.  It  has  long  been  known  that  when 
dilute  solutions  of  two  soluble  salts  are  mixed,  and  no  physical  change 
such  as  the  formation  of  a  solid  or  gas  occurs,  there  is  little  or  no 
heat  change.  Yet  in  many  cases  the  salts  are  capable  of  reacting  on 
each  other,  as  is  shown  by  the  fact  that  on  evaporation  the  reaction 
products  can  be  obtained.  Thus,  when  dilute  solutions  of  sodium 
nitrate  (XaNO3)  and  potassium  chloride  (KC1)  are  mixed,  there  is  no 
heat  change.  On  evaporation  a  certain  quantity  of  potassium  nitrate 


140  GENERAL  CHEMISTRY 

(KNO3)  and  sodium  chloride  (NaCl)  is  obtained,  showing  that  the 
reaction  expressed  in  the  following  equation  may  readily  take. place: 

NaN03  4-  KC1  -  KNO3  +  NaCl 

If  in  dilute  solution  the  various  molecules  represented  in  the  equation 
are  all  parted  into  the  independent  units  K,  Na,  Cl,  NO3,  and  only  in 
concentrated  solution  form  any  considerable  number  of  molecules,  we 
can  see  why  there  should  be  little  evidence  of  reaction  on  mixing  the 
dilute  solutions,  although  in  concentrated  solution  action  takes  place 
accompanied  by  heat  changes. 

4.  The  facts  of  electrolysis.  We  owe  much  of  our  knowledge  as  well 
as  many  of  our  definitions  connected  with  this  important  process  to 
the  Englishman  Michael  Faraday,  who  made  his  largest  contributions 
to  the  subject  in  the  third  decade  of  the  last  century. 

A  substance  whose  solution  will  conduct  the  current  is  termed  an 
electrolyte.  The  electrolyte  is  always  decomposed  during  the  passage 
of  the  current,  and  the  process  of  decomposition  is  called  electrolysis. 

The  plates  dipping  into  the  solution  (Fig.  64) 
are  called  the  electrodes,  the  positive  plate  A, 
by  which  the  current  enters,  being  the  anode, 
the  negative  B,  by  which  it  leaves,  the 
cathode.  The  battery  generating  the  current 
is  represented  by  C.  The  passage  of  the  cur- 
rent is  attended  by  the  movement  of  dis- 
solved particles  through  the  solvent  toward  the  two  electrodes,  and 
it  is  by  means  of  these  that  the  current  is  carried.  The  moving  par- 
ticles are  called  ions,  from  the  Greek  word  meaning  "  to  go,"  those 
moving  toward  the  anode  being  termed  anions,  and  those  moving 
toward  the  cathode,  cations.  During  electrolysis,  decomposition  prod- 
ucts of  the  electrolyte  are  deposited  on  the  electrodes  or  concentrate 
about  them.  For  example,  when  copper  chloride  (CuCl2)  undergoes 
electrolysis,  the  copper  is  deposited  on  the  cathode,  while  the  chlorine 
is  set  free  at  the  anode,  the  products  thus  appearing  at  places  which 
may  be  far  apart. 

Now  Faraday  found  that  when  the  weight  in  grams  of  one  equiva- 
lent of  any  metal  (or  hydrogen)  has  been  liberated  at  the  cathode,  a 
definite  quantity  of  electricity  has  been  carried  through  the  solution, 
namely,  96,540  coulombs.  This  generalization  is  known  as  Faraday's 
law,  and  the  quantity  96,540  coulombs  is  called  a  faraday.  The 


IONIZATION  AND  ELECTROLYSIS  141 

quantity  of  electricity  depends,  therefore,  merely  on  the  number  of 
ions,  not  on  their  character.  Moreover,  under  properly  chosen  con- 
ditions it  can  be  shown  that  very  little  energy  is  absorbed  in  the 
decomposition  of  the  salt.  Now  when  copper  combines  with  gaseous 
chlorine  to  form  copper  chloride,  there  is  a  large  evolution  of  heat: 

Cu  +  C12  =  CuCl2  +  51,630  cal. 

The  decomposition  of  the  salt  into  its  elements  should  therefore  ab- 
sorb the  same  quantity  of  energy.  Yet  when  it  is  dissolved  in  water, 
a  further  evolution  of  heat  takes  place,  equal  to  11,080  cal.,  and 
when  the  resulting  solution  is  electrolyzed  under  the  proper  con- 
ditions, it  is  found  that  there  is  very  little  absorption  of  electrical 
energy  which  it  is  necessary  to  assign  to  the  decomposition  of  the 
salt.  It  would  appear  that  there  is  some  source  of  energy  for  which 
we  have  not  accounted,  which  supplies  the  energy  needful  for  the 
decomposition  of  the  salt. 

All  of  these  facts  —  the  appearance  of  the  products  of  decomposi- 
tion at  distant  electrodes,  the  equal  quantities  of  electricity  which  an 
equivalent  of  every  metal  carries,  the  fact  that  under  properly  chosen 
conditions  there  is  no  absorption  of  energy  due  to  decomposition  of 
the  electrolyte  —  suggest  that  in  the  act  of  solution  there'  is  a  decom- 
position of  the  character  represented  in  the  equilibrium  equation 


and  that  electrolysis  does  not  cause  decomposition  but  depends  upon  it. 
Theory  of  ionization.  With  the  several  laws  relating  to  solutions  of 
electrolytes  and  electrolysis  as  a  foundation,  the  Swedish  chemist 
Arrhenius,  in  1887,  put  forward  a  theory  of  solutions  of  electrolytes 
which  has  come  to  be  known  as  the  theory  of  ionization  or  electrolytic 
dissociation.  The  chief  points  of  this  theory  are  as  follows  : 

1.  Electrolytes  form  ions  in  solution.  When  electrolytes  are  dissolved 
in  water,  their  molecules  tend  to  part  into  two  kinds  of  atoms,  or 
groups  of  atoms,  which  Arrhenius  termed  ions,  following  Faraday's 
usage.    Faraday  formed  no  precise  hypothesis  in  regard  to  the  nature 
of  the  electrical  carriers  in  solution,  but  Arrhenius  made  the  definite 
assumption  that  they  are  not  the  same  as  molecules  but  are  formed 
by  their  decomposition. 

2.  The  ionization  of  the  electrolyte  results  in  an  equilibrium.  The  mole- 
cules of  the  electrolyte  do  not  all  so  break  down,  but  an  equilibrium 


142  GENERAL  CHEMISTRY 

is  presently  reached  between  the  molecules  decomposing  into  ions 
and  those  forming  again  from  the  ions.  In  the  case  of  sodium  chlo- 
ride and  sulfuric  acid,  for  example,  this  may  be  expressed  by  the 
equilibrium  equations 


3.  Various  factors  which  influence  the  equilibrium.   The  point  at  which 
equilibrium  is  reached  depends  largely  upon  three  things  : 

(a)  The  nature  of  the  electrolyte.   Electrolytes   show   great   differ- 
ences among  themselves  as  to  the  extent  to  which  they  give  abnormal 
physical  constants  or  conduct  the  current.    This  is  assumed  to  indicate 
a  varying  percentage  of  ionization. 

(b)  The  nature  of  the  solvent.    This  has  much  influence  upon  the 
ionization,  many  solvents,  such  as  ether  and  benzene,  producing  none 
at  all,  and  others,  such  as  liquid  ammonia  and  sulfur  dioxide,  show- 
ing the  same  effects  as  water,  but  each  to  a  different  degree. 

(c)  The  dilution.    As  the  solution  becomes  more  dilute,  a  larger 
percentage  of  the  solute  is  ionized.    From  a  kinetic  standpoint  this  is 
entirely  reasonable,  for  the  more  widely  the  ions  are  scattered  through 
a  solution  the  less  frequently  they  will  have  opportunity  to  recombine  ; 
so  the  equilibrium  will  move  steadily  toward  a  larger  proportion  of 
ions  as  the  dilution  increases. 

4.  The  ions  are  electrically  charged.    The  theory  must  include  some 
explanation  of  the  way  in  which  the  ions  differ  from  ordinary  mole- 
cules and  atoms.     Ordinary  sodium  and   chlorine  atoms   cannot  be 
present  in  the  equilibrium  expressed  in  the  equation 

NaCl  +=±  Xa  +  Cl 

for  sodium  decomposes  water  violently,  and  chlorine  is  a  gas  but 
slightly  soluble  in  water  and  possessing  a  very  disagreeable  odor, 
whereas  sodium  chloride  forms  an  odorless  and  perfectly  stable  solu- 
tion. Arrhenius  considers  that  these  differences  may  be  accounted  for 
if  we  assume  that  the  ions  are  heavily  charged  with  electricity,  and 
that  in  consequence  of  such  charges  they  have  totally  different  prop- 
erties from  the  uncharged  atoms  and  compounds.  Each  molecule,  on 
ionizing,  yields  two  kinds  of  ions,  having  equal  and  opposite  charges  as 
represented  in  the  equations 


IONIZATION  AND  ELECTROLYSIS  ,    143 

Since  the  opposite  charges  are  always  equal,  the  solution  as  a  whole 
remains  electrically  neutral.  It  is  by  virtue  of  these  charges  on  mov- 
able bodies  that  the  electrolyte  conducts  the  current  through  a  solution. 

Source  of  the  charges  upon  the  ions.  The  present  state  of  advance  in 
our  knowledge  of  electricity  makes  it  possible  to  form  a  more  definite 
conception  of  the.  source  of  the  charges  upon  these  ions.  It  appears  to 
be  very  probable  that  the  atoms  of  the  elements  are  not  homogeneous 
bodies  but  organized  systems,  each  containing  its  own  number  of 
smaller  bodies,  which  are  called  electrons.  These  electrons  are  all 
alike,  and  appear  to  be  about  I^Q-Q  the  weight  of  a  hydrogen  atom. 
They  can  be  separated  from  ordinary  matter,  so  they  are  capable  of 
existing  in  the  free  state.  The  evidence  goes  to  show  that  they  are 
really  negative  electricity,  which  is  therefore  a  material  thing.  A 
body  containing  more  than  its  normal  number  of  electrons  is  said 
to  be  negatively  charged,  while  one  from  which  some  of  its  normal 
number  of  electrons  have  been  removed  is  said  to  be  positively 
charged.  Electrical  energy  is  the  energy  of  innumerable  electrons  in 
very  rapid  motion. 

Applying  these  views  to  the  electrification  of  ions,  we  assume  that 
before  union  the  atoms  of  sodium  and  chlorine  have  each  their  normal 
number  of  electrons.  When  these  combine  to  form  sodium  chloride, 
we  have  no  knowledge  as  to  any  disturbance  in  the  distribution  of  the 
electrons  in  the  several  atoms.  When  the  sodium  chloride  is  dissolved 
in  water,  however,  it  appears  that  this  distribution  tends  to  change. 
The  sodium  atom  loses  one  electron  and  the  chlorine  gains  one. 
The  sodium  atom  is  now  positively  charged,  the  chlorine  negatively, 
and  in  this  condition  they  can  part  from  each  other  to  form  independ- 
ent ions.  Upon  recombination  the  original  condition  of  the  two  atoms 
is  restored,  and  the  molecule  is  electrically  neutral.  The  ions  are 
therefore  very  different  things  from  the  atoms;  they  should  even 
have  different  weights,  though  we  cannot  verify  this  experimentally. 

Application  of  the  theory  of  electrolysis.  The  changes  effected  by 
the  passage  of  the  electric  current  through  solutions  are  in  complete 
accord  with  the  theory  of  ionization,  as  may  be  seen  from  a  study  of 
a  few  typical  examples. 

1.  Electrolysis  of  sodium  chloride.  Experiments  show  that  when 
the  electric  current  is  passed  through  an  aqueous  solution  of  sodium 
chloride  there  are  formed  at  the  cathode  sodium  hydroxide  (XaOH), 
which  remains  in  solution,  and  hydrogen,  which  is  evolved  as  a  gas. 


144     f  GENERAL  CHEMISTRY 

At  the  anode  chlorine  is  set  free  and  either  escapes  in  the  form  of  a 
gas  or  reacts  with  the  constituents  of  the  solution,  according  to  the 
condition  of  the  experiment.  In  terms  of  the  theory  of  ionization  the 
interpretation  of  these  results  is  as  follows : 

Let  Fig.  65  represent  a  solution  of  sodium  chloride,  into  which  the 
electrodes  A  and  B  dip.  The  battery  C  keeps  the  anode  charged  pos- 
itively and  the  cathode  B  negatively.  In  other  words,  it  causes  a  drift 
of  electrons  through  the  wire  from  A  around  to  B,  where  they  accu- 
mulate on  the  cathode,  the  anode 
being  left  deficient  in  them.  Sodium 
chloride  in  solution  ionizes  in  accord- 
ance with  the  equilibrium  equation 

NaCl  ^=>  Na+,  Cl~ 

FlG.  65  r^n  ...  ,  .  '.  XT      4-       • 

The  positive  sodium  ion  Na+  is  at- 
tracted to  the  cathode,  where  it  recovers  its  normal  number  of  elec- 
trons, becomes  an  ordinary  sodium  atom,  and  decomposes  water  as 
follows :  2  Na  +  2  H20  =  2  NaOH  +  H2 

The  chlorine  ion  Cl~,  with  its  excess  electron,  is  attracted  to  the  anode 
A,  to  which  it  gives  up  its  excess,  becomes  ordinary  chlorine,  and 
either  escapes  as  a  gas  or  reacts  with  the  constituents  of  the  solution. 
2.  Electrolysis  of  sulfuric  acid.  When  sulfuric  acid  dissolved  in 
water  is  electrolyzed,  hydrogen  is  evolved  at  the  cathode  and  oxygen 
at  the  anode  (p.  17).  From  the  standpoint  of  the  theory  of  ioniza- 
tion these  changes  result  in  the  following  way :  The  sulfuric  acid  in 
dilute  aqueous  solution  ionizes  as  follows : 

H2S04=<=fcH+,  H+,  SO-  (1) 

The  hydrogen  ions  are  attracted  to  the  negatively  charged  cathode ; 
upon  recovering  their  electrons  from  the  cathode,  they  unite  to  form 
hydrogen  molecules  (H2),  and  escape  as  hydrogen  gas.  The  ions 
SO4~~  are  attracted  to  the  positively  charged  anode,  give  up  their 
excess  electrons,  and  immediately  react  with  the  water  to  form  sul- 
furic acid  and  oxygen,  as  follows : 

42  24  ^    r 

The  oxygen  atoms  unite  to  form  oxygen  molecules  (O2)  and  escape  as 
oxygen  gas.  By  comparing  equations  (1)  and  (2)  it  will  be  noted 
that  the  quantity  of  sulfuric  acid  represented  as  decomposed  in 


IONIZATION  .AND  ELECTKOLYSIS 


145 


equation  (1)  is  the  same  as  that  regenerated  in  equation  (2).  The 
quantity,  therefore,  remains  unchanged,  and  the  process  of  electrolysis 
may  be  continued  as  long  as  any  water  is  presejiit.  It  will  be  noted 
that  the  hydrogen  and  oxygen  evolved  come  indirectly  from  the  water, 
and  are  set  free  in  the  same  ratio  as  that  in  which  they  are  combined 
in  it.  The  part  played  by  the  sulfuric  acid  in  the  electrolysis  of  water 
(p.  11)  is  thus  made  clear. 

3.  Electrolysis  of  sodium  sulfate  (JVa2SOJ.  Sodium  sulf  ate  in  aqueous 
solution  ionizes  as  follows  : 


From  the  standpoint  of  the  theory  of  ionization  one  would  expect  such 
a  solution,  upon  electrolysis,  to  yield  sulfuric  acid  and  oxygen  at  the 
anode  and  sodium  hydroxide  and  hydrogen  at  the  cathode  (see  elec- 
trolysis of  sodium  chloride).  Experiment  shows  that  these  results  are 
actually  obtained.  Upon  electrolysis  the  sul- 
f  ates  of  other  metals  give  similar  results,  form- 
ing sulfuric  acid  and  oxygen  at  the  anode, 
while  at  the  cathode  the  metal  is  either  depos- 
ited as  such  or  reacts  with  the  water  present, 
depending  upon  its  activity  toward  water. 

In  the  electrolysis  of  sodium  sulfate  the  presence 
of  the  base  NaOH  about  the  cathode,  and  the  acid 
H2SO4  about  the  anode,  may  be  demonstrated  as  fol- 
lows :  The  electrolysis  is  effected  in  a  (J-tube  (Fig.  66). 
Before  the  circuit  is  closed,  the  solution  in  the  arm 
of  the  tube  which  contains  the  cathode  is  colored 
with  red  litmus  solution,  while  that  in  the  other  arm 

is  colored  with  blue  litmus.  As  the  electrolysis  proceeds,  it  will  be  found  that  the 
red  litmus  changes  to  blue,  indicating  the  formation  of  a  base,  while  the  blue 
litmus  changes  to  red,  indicating  the  formation  of  an  acid. 

Deductions  from  the  theory  of  ionization.  If  the  theory,  as  outlined, 
presents  a  correct  picture  of  the  conditions  existing  in  a  solution  of 
an  electrolyte,  certain  logical  deductions  at  once  follow: 

1.  Such  a  solution  will  have  two  independent  sets  of  properties, 
the  one  due  to  the  molecules  in  solution,  the  other  to  the  ions.  All 
solutions  containing  a  certain  ion,  say  Cl~,  should  have  one  set  of 
properties,  irrespective  of  the  source  of  that  ion.  This  is  found  to  be 
true.  For  example,  all  chlorides  which  are  electrolytes,  when  treated 
with  a  solution  of  silver  nitrate  (AgNO3),  precipitate  insoluble  silver 


FIG.  66 


146  GENERAL  CHEMISTRY 

chloride  (AgCl).  This  may  be  regarded  as  due  to  an  equilibrium  which 
the  silver  and  chlorine  ions  tend  to  set  up  with  silver  chloride  : 


But  since  the  latter  salt  is  practically  insoluble,  as  soon  as  a  very 
little  of  it  is  formed  the  solution  becomes  supersaturated  and  a 
precipitate  results. 

2.  Since  in  very  dilute  solution  we  assume  that  the  ionization  is 
nearly  complete,  the  color  of  the  solution  must  be  due  to  the  several 
ions,  rather  than  to  the  molecules,  while  in  more  concentrated  solution 
the  color  due  to  the  molecules  will  predominate.  In  accordance  with  this 
view  we  find  that  concentrated  solutions  of  various  salts  of  the  same 
colored  metal  have  quite  a  variety  of  colors,  while  then*  dilute  solu- 
tions have  the  same  color.    Thus,  concentrated  solutions  of  copper 
salts  have  various  shades  of  blue,  yellow,  and  green,  while  their  dilute 
solutions,  which  give  the  copper  ion  Cu++,  are  all  pale  blue  and  are 
not  distinguishable  in  color. 

3.  In  very  dilute  solutions  in  which  the  ionization  is  regarded  as 
practically  complete,  the  effect  of  the  dissolved  electrolyte  upon  the 
boiling  and  freezing  points  of  the  solvent,  as  well  as  upon  the  con- 
ductivity and  osmotic  pressure  of  the  solution,  should  reach  a  maxi- 
mum value.    Thus,  the  effect  produced  by  sodium  chloride  (NaCl)  on 
complete  ionization  should  be  limited  to  twice  that  produced  by  an 
equivalent  quantity  of  a  nonelectrolyte,  such  as  sugar,  while  barium 
chloride  (BaCl2),  which  forms  three  ions,  should  have  a  maximum 
effect  of  three  times  that  of  sugar.    In  a  general  way  it  may  be  said 
that  such  is  the  fact,  though  quantitative  experiments  do  not  always 
give  the  results  which  would  be  expected. 

The  quantitative  estimation  of  ionization.  A  discussion  of  the  quantitative 
estimation  of  the  extent  of  ionization  would  take  us  too  far  for  the  scope  of 
this  book.  Such  measurements  are  based  upon  the  extent  to  which  solutions 
of  electrolytes  deviate  from  those  of  nonelectrolytes,  and  upon  measurements  of 
the  electrical  conductivity  of  solutions.  It  must  be  said  that  these  methods  do 
not  always  yield  satisfactory  agreement  among  themselves.  When  we  remember 
the  widely  different  conditions  under  which  the  experiments  are  made,  —  at  the 
freezing  point,  at  the  boiling  point,  and  at  ordinary  room  temperature,  —  the  lack 
of  agreement  is  not  surprising.  Moreover,  we  know  very  little  as  to  the  extent  to 
which  chemical  combination  between  solvent  and  solute  enters  in  to  complicate 
the  matter.  In  general,  there  are  so  many  complicating  circumstances  of  which 
we  are  aware,  and,  doubtless,  so  many  others  of  which  we  know  nothing,  that  it 
may  be  said  the  methods  agree  as  closely  as  could  be  expected. 


IONIZATION  AND  ELECTKOLYSIS  147 

Summary.  In  the  theory  of  ionization  we  have  to  do  with  another 
broad  theory,  which  endeavors  to  give  us  a  working  picture  of  the 
peculiarities  of  the  class  of  substances  known  as  electrolytes.  This 
picture  has  already  been  considerably  modified  since  it  was  first  drawn, 
and  will,  no  doubt,  be  further  modified  in  the  future.  At  many  points 
it  is  unsatisfactory,  for  it  has  to  do  with  one  of  the  most  complicated 
provinces  of  chemistry.  It  has,  however,  stimulated  a  vast  amount  of 
research,  has  greatly  extended  our  knowledge,  and  at  the  present 
tune  gives  us  by  far  the  most  satisfactory  conception  of  solutions  that 
we  have.  A  great  many  of  the  facts  of  chemistry  can  be  presented 
much  more  simply  in  terms  of  this  theory  than  in  any  other  way  open 
to  us  at  present.  If  we  keep  before  us  the  limitations  of  all  theory, 
we  shall  be  able  to  use  this  one  to  great  advantage. 


CHAPTER  XIII 

NEUTRALIZATION 

Introduction.  The  great  majority  of  the  compounds  to  be  described 
in  the  course  of  our  study  belong  to  one  of  three  classes,  namely, 
acids,  bases,  and  salts,  to  which  reference  has  been  repeatedly  made 
in  the  preceding  pages.  The  individual  substances  will  be  described  in 
detail  in  subsequent  chapters,  but  it  is  important,  before  proceeding 
farther,  to  become  familiar  with  their  properties  as  classes. 

Acids.  An  inspection  of  the  formulas  of  the  acids  shows  that  they 
all  contain  hydrogen,  but  many  other  substances,  such  as  sugars,  oils, 
fats,  and  waxes,  also  contain  hydrogen  and  yet  are  not  acids.  As  a 
class  acids  are  characterized  by  their  ability  to  change  the  colors  of 
certain  organic  compounds.  For  example,  they  turn  blue  litmus  red, 
while  red  phenolphthalein  is  rendered  colorless.  Those  which  are 
appreciably  soluble  usually  have  a  more  or  less  decidedly  sour  taste, 
but  some  are  too  insoluble  to  exhibit  this  property,  and  in  some  sol- 
uble ones  it  is  not  at  all  noticeable.  They  all  act  chemically  upon 
the  class  of  bodies  known  as  bases,  in  which  action  then-  own  dis- 
tinctive properties  are  of  course  lost.  In  dilute  aqueous  solution  they 
are  conductors  of  the  electric  current.  If  we  make  use  of  the  con- 
ceptions of  the  ionization  theory,  we  can  say  that  they  all  yield 
hydrogen  ions.  An  acid  may  then  be  denned  as  any  compound  which 
produces  hydrogen  ions  when  dissolved  in  water.  This  gives  us  our 
best  definition  of  an  acid,  for  it  is  to  the  presence  of  the  hydrogen 
ion,  and  not  to  any  property  possessed  by  the  molecule  of  the  com- 
pound, that  the  characteristic  conduct  of  an  acid  is  attributed. 

1.  Characteristics  of  the  hydrogen  ion.  The  most  prominent  character- 
istic of  the  hydrogen  ion  is  its  tendency  to  combine  with  the  hydroxyl 
ion  OH~,  forming  water  (H2O).  Water  is  ionized  to  but  a  very  slight 
extent,  so  that  when  a  dilute  solution  of  an  acid  is  brought  into  contact 
with  any  substance  which  forms  hydroxyl  ions,  the  hydrogen  ions,  and 
therefore  all  the  characteristic  properties  of  the  acid,  disappear.  An 
example  of  such  action  is  represented  in  the  equation 

K+,  OH-  +  H+,  Cl-  =  H2O  +  K+,  Cl- 

148 


KEUTBALIZATION  149 

A  second  characteristic  property  of  the  hydrogen  ion  is  its  conduct 
toward  many  metals.  When  metals  such  as  magnesium,  zinc,  and 
iron  are  exposed  to  the  action  of  dilute  acids,  the  metallic  atoms  pass 
into  solution  as  ions,  while  an  equivalent  weight  of  hydrogen  ions 
lose  their  charge  and  escape  as  gas.  An  example  of  this  is  repre- 
sented in  the  equation 


Zn  +  2(H+,  Cl-)=  Zn++,  2C1-  +  H 


This  action  may  be  described  in  terms  of  electrons  by  saying  that  the 
hydrogen  ion  has  a  stronger  affinity  for  an  electron  than  has  the  zinc 
atom.  Consequently,  the  hydrogen  ion  recovers  its  missing  electron 
from  the  zinc,  which  in  turn  becomes  a  positive  ion  and  passes  into 
solution. 

2.  Basicity  of  acids.  The  formulas  of  the  three  acids,  hydrochloric 
(HC1),  sulfuric  (H2SO4),  and  phosphoric  (HgPO4),  show  that  in  one 
molecule  there  may  be  several  hydrogen  atoms.    Experiment  shows 
that  all  of  them  may  become  ions  and  so  suffer  replacement  by  metals. 
The  number  of  ionic  or  replaceable  hydrogen  atoms  in  a  molecule 
of  an  acid  is  called  the  basicity  of  the  acid.    Hydrochloric  acid  is 
monobasic,  sulfuric  is  dibasic,  and  phosphoric  is  tribasic.     The  for- 
mulas of  some  acids,  such  as  acetic  (C2H4O2)  and  tartaric  (C4HgO6), 
would  apparently  indicate  a  still  higher  basicity,   but    experiment 
shows  that  in  the  former  only  one,  and  in  the  latter  two,  hydrogen 
atoms  act  as  ions.    On  this  account  the  formulas  are  usually  written 
H  •  C0HgO2  and  H2  •  C4H4Og.   Groups  of  atoms  which  go  through  reac- 
tions undisturbed,  acting  as  a  unit,  like  the  group  C2HgO2  of  acetic 
acid  or  the  group  SO4  of  sulfuric  acid,  especially  when  they  can  play 
the  role  of  an  ion,  are  called  radicals. 

3.  Strength  of  acids.    Even  a  slight  acquaintance  with  acids  will 
convince  one  that  they  differ  greatly  in  strength.     Since  the   acid 
properties  are  attributed  to  the  hydrogen  ion,  that  acid  should  be  the 
strongest  which,  for  a  given  weight  of  hydrogen  in  its  solution,  pro- 
duces the  most  hydrogen  ions.    This  cannot  be  ascertained  from  the 
formula,  for  ionization  is  always  an  equilibrium,  and  the  concentration 
of  the  hydrogen  ions  depends  not  only  upon  the  weight  of  hydrogen 
in  the  acid  of  a  given  solution,  but  also  upon  the  percentage  of  the 
molecules  which  are  ionized  when  equilibrium  is  reached.   In  the  table 
on  page  155  the  strength  of  the  common  acids  is  expressed  in  per- 
centage of  ionization  under  definite  conditions. 


150  GENERAL  CHEMISTRY 

4.  The  ionization  of  dibasic  and  tribasic  acids.  In  the  case  of  dibasic 
acids  such  as  sulfuric  acid  (H2SO4),  experiment  shows  that  in  mod- 
erately concentrated  solutions  the  ionization  is  largely  as  follows  : 


In  quite  dilute  solutions,  however,  a  second  stage  of  ionization  is 
reached  :  -  __>.  H+  +  $O-- 


Jn  more  concentrated  solutions,  therefore,  sulfuric  acid  gives  a  much 
smaller  percentage  of  its  hydrogen  in  the  form  of  ions  than  does  a 
strong  monobasic  acid  such  as  hydrochloric  acid,  and  the  table  shows 
that  it  is  rated  as  much  weaker.  With  the  tribasic  phosphoric  acid 
H3PO4  the  ionization  takes  place  in  three  successive  stages,  with 
increasing  dilution,  and  may  be  represented  as  follows  : 

H3P04  +=±  H+  +  H2P04-  +=+  2  H+  +  HP04~  ^=±  3  H+  +  PO4~ 

The  ionization  expressed  in  the  final  stage  is  very  slight. 

5.  Nonioniziny  solvents.  There  are  many  solvents  in  which  at  least 
some  of  the  acids  are  soluble,  yet  which  do  not  occasion  ionization  of 
the  acid.  In  tne  absence  of  all  traces  of  water  such  solutions  do  not 
have  the  properties  which  we  associate  with  the  presence  of  acids.  It 
should  be  said,  however,  that  such  solutions  present  many  features 
not  easily  interpreted  in  terms  of  the  ionization  theory. 

Bases,  or  metallic  hydroxides.  In  composition,  the  bases  are  made  up 
of  some  metallic  element  (or  radical)  in  combination  with  oxygen  and 
hydrogen.  The  strongest  and  most  soluble  representatives  of  the  group, 
namely,  sodium  hydroxide  (NaOH)  and  potassium  hydroxide  (KOH), 
are  called  alkalies.  Many  of  the  bases  are  practically  insoluble  in  water. 
When  soluble,  they  reverse  the  color  changes  produced  in  indicators 
by  acids,  turning  red  litmus  blue  and  colorless  phenolphthalein  red. 
The  soluble  ones  have  a  more  or  less  brackish  taste  and  feel  like  soap. 
The  following  table  gives  the  color  changes  of  a  few  more  commonly 
employed  indicators. 

TABLE  OF  INDICATORS 


NAME  OF  INDICATOR 

ACID  SOLUTION 

BASIC  SOLUTION 

Litmus  .... 

Red 

Blue 

Phenolphthalein       .           .     . 
Methyl  orange    
Congo  red  . 

Colorless 
Red 
Blue 

Purplish-red 
Yellow 
Red 

Cochineal  . 

Yellowish-red 

Purple 

NEUTRALIZATION  151 

In  solution  the  bases  act  chemically  upon  the  acids,  and  both  lose 
their  characteristic  properties.  This  reaction  may  be  represented  in 
the  following  way:  All  bases  furnish  hydroxyl  ions  (OH~)  in  solu- 
tion, as  indicated  in  the  equations 


Ca  (OH)2  +=±  Ca+  +  2  (OH~) 
BiO-OH+=tBiO++OH- 

In  the  presence  of  an  acid  the  hydroxyl  ion  of  the  base  combines 
with  the  hydrogen  ion  of  the  acid  to  form  water  : 

H+  +  OH-  =  H20 

This  reaction  removes  the  characteristic  ions  of  both  acid  and  base 
from  the  solution.  A  base  may  therefore  be  denned  as  a  compound 
which  produces  hydroxyl  ions  when  dissolved  in  water. 

1.  Strength  of  bases.    The  ionization  of  a  base,  like  that  of  an  acid, 
is  a  reversible  reaction  leading  to  an  equilibrium,  and  the  percentage 
of  ionization  when  equilibrium  is  reached  varies  greatly  with  differ- 
ent bases.    As  with  acids,  the  largely  ionized  bases  are  the  strong  ones, 
while  those  which  are  little  ionized  are  weak.    The  table  on  page  155 
gives  the  percentage  of  ionization  of  some  of  the  commonest  bases 
under  stated  conditions. 

2.  Acidity  of  bases.  The  formulas  for  potassium  hydroxide  (KOH), 
calcium  hydroxide  (Ca(OH)2),  and  aluminium  hydroxide  (Al  (OH)3) 
show  that  molecules  of  bases  may  contain  several  hydroxyl  groups. 
When  the  molecule  can  furnish  but  one  hydroxyl  ion,  the  base  is  said  to 
be  a  monacid  base  ;  when  two,  a  diacid  base  ;  when  three,  a  triacid  base. 
There  are  a  few  Utracid  bases,  such  as  stannic  hydroxide  (Sn(OH)4), 
but  for  the  most  part  these  lose  a  molecule  of  water  and  act  as  diacid 

Sn(OH)4=  SnO(OH)2  +  H2O 

The  diacid  and  triacid  bases,  like  the  dibasic  and  tribasic  acids,  ionize 
in  stages,  so  that  for  a  given  weight  of  hydroxyl  in  solution  they  are 
not  as  strong  as  are  the  monacid  bases.  Few  of  these  bases  are  soluble 
enough  to  permit  of  the  ready  comparisons  which  may  be  easily  made 
in  the  case  of  most  of  the  acids. 

Salts.  When  an  acid  and  a  base  are  brought  together  in  solution,  the 
reaction  always  consists  in  the  union  of  the  hydrogen  and  hydroxyl  ions 
to  form  water  : 

K+,  OH-  +  H+,  Cl-  =  H2O  +  K+,  Cl- 


152  GENERAL  CHEMISTEY 

This  action  is  called  neutralization.  The  remaining  positive  and  nega- 
tive ions,  whose  electrical  charges  always  balance  each  other,  then  set 
up  their  own  equilibrium,  as  represented  in  the  following  equation : 

K++  Cl-  +=*  KC1 

When  the  solution  becomes  saturated  with  the  molecular  member  of 
the  equilibrium  (in  this  case  KC1),  either  because  of  its  limited  solu- 
bility or  through  evaporation  of  the  solvent,  the  excess  separates 
in  solid  form  and  is  known  as  a  salt.  A  salt  may  therefore  be  defined 
as  a  compound  composed  of  the  anion  of  any  acid  and  the  cation  of 
any  base.  Since  the  cations  of  the  inorganic  bases  are  metallic  ele- 
ments, we  may  regard  a  salt,  apart  from  any  theoretical  considerations, 
as  an  acid  in  which  the  hydrogen  has  been  replaced  by  a  metal.  As 
we  have  seen,  they  may  sometimes  be  prepared  directly  by  the  action 
of  an  acid  upon  a  metal. 

Hydrogen  salts.  If  we  regard  acid  properties  as  due  solely  to  the 
presence  of  hydrogen  ions,  which  exist  only  in  solution,  it  is  evident 
that  when  we  obtain  an  acid  as  a  crystalline  solid,  such  as  oxalic  acid 
(H2C2O4  •  2H2O),  or  in  pure  liquid  form,  as  nitric  acid  (HNOg),  or  as 
a  dry  gas  (HC1),  these  acid  properties  should  be  absent.  This  is  the 
case,  such  substances  being  without  action  on  indicators.  It  is  well, 
therefore,  to  regard  the  pure  substances  as  salts  of  hydrogen,  reserv- 
ing the  term  acid  for  their  solutions  in  ionizing  solvents.  Thus,  the 
gas  HC1  is  called  hydrogen  chloride ;  the  pure  liquid  HNO3,  hydro- 
gen nitrate.  This  usage  will  be  frequently  followed  in  this  text, 
though  it  is  not  always  convenient  to  observe  it  rigidly. 

General  properties  of  salts.  Since  there  is  no  one  ion  characteristic 
of  salts,  there  is  no  set  of  properties  which  they  all  have  in  common, 
save  that  they  are  electrolytes.  In  solubility  they  range  between  the 
widest  limits,  some  dissolving  to  an  extent  of  less  than  a  milligram 
per  liter,  and  others  dissolving  in  much  less  than  their  own  weight  of 
water.  In  degree  of  ionization  the  salts  are  found  to  be  much  more 
nearly  the  same  than  is  the  case  with  either  acids  or  bases.  In  a  general 
way  it  may  be  said  that  they  are  all  strongly  ionized  to  about  the 
same  extent  as  strong  acids  and  bases.  The  salts  of  magnesium,  cad- 
mium, zinc,  and  mercury  are  exceptional,  being  much  less  ionized  than 
most  other  salts.  The  table  on  page  155  gives  values  for  a  few  rep- 
resentative salts.  Nearly  all  the  salts  we  shall  meet  with  are  solids, 
and  the  great  majority  are  colorless.  The  color  of  a  salt  is  due  in  part 


NEUTRALIZATION 


153 


to  the  acid  and  in  part  to  the  base  from  which  the  salt  is  derived. 
When  both  are  colorless  (as  is  true  with  the  majority  of  both  acids 
and  bases),  the  salt  is  likely  to  be  colorless,  or  nearly  so.  When  either 
has  a  marked  color,  the  salts  are  all  apt  to  share  it  in  some  degree. 
Thus,  copper  salts  are  usually  some  shade  of  blue  or  green,  while  the 
salts  of  the  reddish  chromic  acid  are  usually  yellow  or  orange. 

Preparation  of  salts  from  oxides.  Nearly  all  of  the  metallic  hydrox- 
ides, when  heated,  readily  form  the  corresponding  oxide.  With  calcium 
hydroxide  the  equation  is  as  follows  : 


A  number  of  oxides,  among  others  that  of  calcium,  are  known  to 
combine  with  water  to  form  the  hydroxide,  so  it  seems  probable  that 
there  is  always  an  equilibrium  between  an  oxide,  water,  and  the  corre- 
sponding hydroxide.  It  is  not  surprising,  therefore,  that  oxides  sus- 
pended in  water  act  upon  acids  just  as  do  hydroxides  to  form  salts. 
Thus,  copper  oxide  dissolves  in  hydrochloric  acid  to  form  copper 
chloride  : 


H20 


CuO  +  2  HC1  =  CuCl2 

In  a  great  many  cases  this  is  the  most  con- 
venient way  in  which  to  prepare  salts,  since 
most  of  the  oxides  are  more  easily  obtained 
than  the  corresponding  hydroxides. 

Neutralization  a  definite  act.  If  two  solu- 
tions, one  of  a  base  and  the  other  of  an  acid, 
are  prepared,  experiment  has  shown  that  a 
given  volume  of  the  acid  will  invariably 
require  a  perfectly  definite  volume  of  the 
base  for  its  neutralization.  The  experiment 
is  most  easily  performed  with  the  aid  of 
burettes  (Fig  67),  which  are  graduated 
tubes  furnished  with  a  stopcock  at  the  end. 
The  one  is  filled  to  the  zero  mark  with  the 
acid  solution,  the  other  with  the  basic.  A 
measured  volume  of  the  one  solution  is  drawn  off  into  a  small  beaker, 
a  few  drops  of  an  appropriate  indicator  added,  and  the  second  solution 
run  in  with  constant  stirring  until  the  indicator  just  turns  color.  The 
process  just  described  is  called  titration.  If  the  concentration  of  each 
solution  is  accurately  known,  it  is  easy  to  calculate,  from  the  volumes 


t 


FIG. 67 


154  GENERAL  CHEMISTRY 

required  for  neutralization,  the  ratio  by  weight  between  the  acid  and 
the  base  taking  part  in  the  action.  Experiment  shows  that  this  ratio 
always  bears  a  simple  relation  to  that  between  the  molecular  weights 
of  the  reacting  substances.  Such  a  reaction  as  is  indicated  in  the 
equation  -TT+  rur- _L  TI+  ni-  TT+  r»i-  _i_  TT  n 

JV   ,  v/rl    -f-  ±1   ,  \^il    =  Jv   ,  \ji    -f-  Jtl  \J 

is  therefore  perfectly  definite,  and  does  not  in  general  stop  short 
of  completion. 

Reference  to  the  table  on  page  155  will  show  that  the  solution  of 
ammonium  hydroxide  is  dissociated  to  the  extent  of  only  0.3  per  cent. 
It  might  be  thought  from  this  that  when  such  a  solution  is  titrated  with 
an  acid,  complete  neutralization  will  occur  when  0.3  per  cent  of  the 
base  has  been  acted  on  by  the  acid.  This  is  not  the  case,  the  indicator 
changing  color  only  when  all  of  the  base  has  been  acted  upon.  In 
explanation,  it  must  be  remembered  that  the  condition  of  ammonium 
hydroxide  in  solution  is  an  equilibrium : 

NH4OH  <=±  NH4+  +  OH- 

As  fast  as  the  hydroxyl  ions  (OH~)  are  removed  by  the  reaction  of 
neutralization,  more  are  supplied  by  ionization  of  the  base  NH4OH, 
so  that  all  of  the  base  is  finally  brought  into  reaction. 

Normal  solutions.  In  scientific  investigation,  as  well  as  in  industrial 
analysis,  it  is  often  desirable  to  estimate  the  weight  of  acid  or  base  in 
a  given  volume  of  solution.  For  example,  the  acid  in  vinegar  and  the 
alkali  in  lye  or  limewater  must  be  determined  very  frequently.  This 
may  be  readily  accomplished  by  titrating  the  solution  of  unknown  con- 
centration with  a  solution  of  an  acid  or  a  base  the  concentration  of 
which  is  accurately  known.  Such  a  solution  is  called  a  standard  solu- 
tion. The  standard  may  be  of  any  convenient  concentration,  depend- 
ing upon  the  character  of  the  solutions  to  be  investigated.  If  the 
standard  solution  is  prepared  in  such  a  way  that  1  1.  will  contain  1  g. 
equivalent  of  hydrogen  (1.008  g.  H)  or  of  hydroxyl  (17.008  g.  OH), 
then  the  solution  is  called  a  normal  solution.  Solutions  of  half  this 
concentration  are  half  normal  (N/2).  Tenth-normal  solutions  (N/10) 
are  still  more  frequently  employed. 

Such  solutions  greatly  simplify  calculations,  for  it  is  evident  that  1 1.  of  a 
normal  solution  of  any  acid  will  neutralize  1  molar  weight  of  any  monacid  base ; 
10  cc.  will  neutralize  -jj^  of  its  molar  weight.  If  in  titrating  a  solution  of  lye 
(NaOH)  of  unknown  concentration  20  cc.  of  normal  acid  is  required,  then  the 
solution  of  lye  contains  Tf$7  x  40.01  (=  0.8002)  g.  NaOH. 

v  -"'••'  ' 

\  CrcrO  > 

V  O 


NEUTRALIZATION 


155 


Relative  ionization  of  electrolytes.  Since  normal  solutions  are  all 
chemically  equivalent  to  each  other,  it  is  convenient  to  indicate  the 
degree  of  ionization  of  an  electrolyte  by  stating  the  percentage  ionized 
in  its  normal  solution.  The  following  tables  will  be  found  convenient 
for  reference,  though  the  values  are  merely  approximate : 

TABLE    OF  IONIZATION  IN  NORMAL   SOLUTIONS 


ACIDS  :  Per  cent 

Nitric 85.0 

Hydrochloric 79.0 

Sulfuric 51.0 

Acetic 0.3 

BASES : 

Potassium  hydroxide       .     .  76.0 

Sodium  hydroxide      .     .     .  72.0 

Barium  hydroxide       .     .     .  69.0 

Ammonium  hydroxide    .     .  0.3 

Calcium  hydroxide  (N/64)  .  90.0 
(Veiy  little  soluble) 


SALTS  :  Per  cent 

Potassium  chloride     .     .     .  75.5 

Sodium  chloride     ....  68.6 

Ammonium  chloride  .     .     .  74.0 

Potassium  nitrate  ....  63.5 

Potassium  acetate       .     .     .  63.5 

Silver  nitrate 58.5 

Potassium  sulfate  ....  54.7 

Potassium  carbonate  .     .     .  49.0 

Sodium  sulfate 46.0 

Zinc  sulfate 23.0 

Copper  sulfate  .  .     .  22.5 

Mercury  salts     .     .     .    Very  little 


TABLE   OF   IONIZATION   AT   OTHER   CONCENTRATIONS 


ACIDS  IN  N/10  SOLUTION  : 

Per  cent 

Acetic 1.1 

Carbonic 0.17 

Hydrosulfuric 0.07 

Hydrocyanic 0.01 


ACIDS  IN  CONCENTRATED  SOLUTION  : 

Per  cent 
Hydrochloric  (35%)    .     .     .     13.6 

Nitric  (62%) 9.6 

Sulfuric  (95%)  .....       0.7 


Varieties  of  salts.  It  is  evident  that  hydrochloric  acid  (HC1)  and 
potassium  hydroxide  (KOH)  can  act  upon  each  other  in  but  one 
proportion,  namely,  that  required  by  the  following  equation : 

HC1  +  KOH  =  KC1  +  H2O 

With  the  base  copper  hydroxide  (Cu(OH)2)  there  are  two  possibilities: 

Cu(OH)2  +  2  HC1  =  CuCl2  +  2  H2O 
Cu(OH)2  +  HC1  =  Cu(OH)Cl  +  H2O 

In  the  second  instance  only  one  of  the  two  hydroxyl  groups  of  the 
base  has  entered  into  reaction  with  the  acid.  The  substance  Cu(OH)Cl 


156  GENERAL   CHEMISTRY 

is  evidently  a  salt,  for  it  contains  copper  and  chlorine,  both  capable  of 
forming  ions  in  solution.  It  is  also  a  base,  for  it  contains  the  radical 
OH,  which  readily  forms  the  ion  OH~  in  solution.  Since  it  combines 
both  of  these  characteristics,  it  is  called  a  basic  salt.  A  basic  salt  may 
therefore  be  defined  as  a  base  partially  neutralized  by  an  acid. 

Similarly,  with  sulfuric  acid  (H2SO4),  potassium  hydroxide  may 
have  two  possible  reactions: 

H2SO4  +  2  KOH  =  K2SO4  +  2  H2O 
"  H2S04  +  KOH  =  KHS04  +  H2O 

The  substance  KHSO4,  called  potassium  hydrogen  sulfate,  or  potassium 
acid  sulfate,  yields  in  solution  the  ions  K+,  H+,  SO4~~.  Of  these, 
K+  and  SO4~~  are  typical  salt  ions,  while  the  other,  H+,  is  the  charac- 
teristic ion  of  acids.  The  substance  KHSO4  is  therefore  called  an  acid 
salt,  which  may  be  defined  as  an  acid  partially  neutralized  by  a  base. 
Salts  of  the  usual  type,  such  as  sodium  chloride  (NaCl)  or  sulfate 
(Na2SO4),  which  form  neither  hydroxyl  nor  hydrogen  ions  in  solution, 
but  only  the  anion  of  an  acid  and  the  metallic  cation  of  a  base,  are 
called  normal  salts. 

Mixed  salts.  We  may  also  have  salts  of  the  type  KNaSO4  or  Ca(NO8)Cl,  in 
which  one  anion  combines  with  two  different  metallic  cations,  or  one  cation  with 
two  different  anions.  Such  salts  are  called  mixed  salts.  The  former  type  is  far  the 
more  common,  and  we  shall  have  occasional  examples  of  such  mixed  types  as  we 
proceed. 

Preparation  of  acid  and  basic  salts.  A  few  words  in  regard  to  the 
preparation  of  acid  and  basic  salts  will  serve  to  emphasize  their 
characteristics. 

Basic  salts  may  be  prepared  in  either  of  two  general  ways  :  (1)  by 
the  partial  neutralization  of  a  base  (which  must  have  an  acidity  of 
more  than  1)  : 

Mg(OH)2  +  HC1  =  Mg(OH)Cl  +  H2O 

(2)  by  the  action  of  a  base  on  some  normal  salt  of  the  same  base  : 
Mg(OH)2  +  MgCl2  =  2  Mg(OH)Cl 

Acid  salts  may  be  prepared  by  analogous  reactions:  (1)  by  the 
partial  neutralization  of  an  acid  (which  must  have  a  basicity  of  more 
than  1)  : 


(2)  by  the  action  of  such  an  acid  on  a  normal  salt  of  the  same  acid: 


Na2S04  +  H2S04  =  2  NaHSO4 


NEUTRALIZATION  157 

The  thermochemistry  of  neutralization.  Since  the  heat  of  reaction 
affords  an  insight  into  the  intensity  of  the  attraction  between  sub- 
stances, experiments  have  been  made  by  many  chemists  on  the  heat 
of  reaction  when  various  bases  act  upon  acids.  The  Danish  chemist 
Julius  Thomsen  was  especially  prominent  in  such  thermal  deter- 
minations, and  his  results  show  that  when  the  neutralization  is  con- 
ducted in  dilute  solution,  and  involves  a  strong  acid  and  base,  the 
heat  of  reaction  is  approximately  the  same  in  every  case,  the  average 
value  being  about  13,790  cal.  The  following  equations  show  some  of 
Thomson's  results. 

HEAT  OF  NEUTRALIZATION:  STRONG  ELECTROLYTES 


NaOH 

+ 

HC1 

= 

NaCl 

+ 

11 

2o 

+ 

13,780  cal. 

NaOH 

+ 

HNO, 

= 

NaNO3 

+ 

II 

2o 

+ 

13,680  cal. 

NaOH 

+ 

HBr 

= 

NaBr 

+ 

II 

2o 

+ 

13,750  cal. 

NaOH 

+ 

HI03 

= 

NaIO3 

+ 

II 

20 

+ 

13,810  cal. 

HC1 

+ 

KOH 

= 

KC1 

+ 

II 

2o 

+ 

13,750  cal. 

2  HC1   +  Ba(OH)2  =  BaCl2     +  2  H2O  +  27,780  cal. 
2  HC1  +  Ca(OH)2  =  CaCl2     +  2  H2O  +  27,900  cal. 

This  unexpected  result  is  in  close  accord  with  the  theory  of  ionization, 
which  assumes  strong  acids  and  bases  to  be  largely  ionized  in  solu- 
tions such  as  Thomsen  used  (1  :  200).  In  all  such  cases  the  reaction 
taking  place  is  the  same,  namely,  the  union  of  the  hydrogen  with  the 
hydroxyl  ion  : 

H+  +  OH-  =  H2O  +  13,790  cal. 

In  the  case  of  weak  acids  or  bases  the  ionization  is  only  partial, 
even  in  dilute  solutions.  As  neutralization  proceeds  more  molecules 
are  changed  into  ions.  This  change  is  a  true  chemical  reaction  and  is 
attended  by  a  heat  change  which  may  be  either  positive  or  negative. 
Accordingly,  with  weak  acids  or  bases  the  heat  of  neutralization  may 
be  either  greater  or  less  than  13,790,  and  Thomsen  found  examples 
of  both  kinds. 

HEAT  OF  NEUTRALIZATION:  WEAK  ELECTROLYTES 

HC1  +  NH4OH  =  NH4C1  +  H2O  +  12,270  cal. 
2  HC1  +  Fe(OH)2  =  FeCl2  +  2  H2O  +  21,390  cal. 
2  HC1  +  Zn(OH)2  =  ZnCl2  +  2  H2O  +  19,880  cal. 
NaOH  +  HF  =  NaF  +  H2O  +  16,272  cal. 

NaOH  +  HPO3  =  NaPO3  +  H2O  +  14,380  cal. 
NaOH  +  HC1O  =  NaCIO  +  H2O  +  9,980  cal. 


158 


GENERAL   CHEMISTRY 


The  electromotive  series  ;  electrode  potential.  It  has  been  stated  that 
when  certain  metals  are  brought  into  contact  with  an  acid  solution, 
hydrogen  ions  are  discharged,  an  equivalent  weight  of  the  metal 
passing  into  solution  in  the  ionic  condition.  This  reaction  is  attended 
by  considerable  heat  evolution,  and  experiment  shows  that  the  heat  of 
reaction  with  various  metals  is  quite  different.  For  example,  Thomsen 
found  the  following  values  for  magnesium,  zinc,  and  iron : 

Mg  +  2  (H+,  C1-)  -  Mg+  +,  2  Cl-  +  H2  +  108,290  cal. 
Zn  +  2  (H+,  C1-)  =  Zn++,  2  Cl~  +  H2  +  34,210  cal. 
Fe  +  2  (H+,  C1-)  =  Fe+  +,  2  Cl-  +  H2  +  21,320  cal. 

Now  108,290  -  34,210  =  74,080  cal.  represents  the  difference  m  the 
energy  liberated  in  the  solution  of  a  gram  atomic  weight  of  magnesium 
as  compared  with  one  of  zinc.  If,  therefore;  we  bring  magnesium  into 
a  solution  of  a  ziric  salt,  we  should  expect  the  zinc  to  be  displaced  by 
the  magnesium,  with  the  liberation  of  74,080  cal.  This  is  found  to  be 
the  case.  A  very  convenient  experiment  of  the  same  kind  can  be 
made  by  sprinkling  zinc  dust  into  a  solution  of  copper  sulfate.  The 
blue  color  of  the  solution  soon  disappears,,  and  dark-brown  copper  pow- 
der is  precipitated  in  place  of  the  gray  zinc.  A  thermometer  in  the 
solution  will  show  a  considerable  rise  in  temperature  at  the  same  time. 
It  is  possible  to  arrange  all  the  metals  in  the  order  in  which  they  will 
displace  each  other  in  this  way.  Such  an  arrangement  is  known  as  the 
electromotive  series  of  the  metals  and  is  given  in  the  following  table : 


ELECTROMOTIVE 

SERIES 

1. 

Caesium 

8. 

Aluminium 

15. 

Nickel 

22. 

Bismuth 

2. 

ft. 

Rubidium 
Potassium 

9. 
10. 

Manganese 
Zinc 

16. 
17. 

Tin 
Lead 

23. 
24. 

Mercury 

Silver 

4. 
5. 

Sodium 
Lithium 

11. 
12. 

Chromium 
Cadmium 

18. 
19. 

Hydrogen 
Arsenic 

25. 

2G. 

Platinum 
Gold 

0. 
7. 

Calcium 
Magnesium 

13. 
14. 

Iron 
Cobalt 

20. 
21. 

Copper 
Antimony 

All  the  elements  above  a  given  one  in  the  series,  or,  as  it  is  usually 
expressed,  all  those  having  a  higher  electrode  potential,  will  displace  it 
from  solution,  while  it  in  turn  will  displace  all  those  of  lower  electrode 
potential.  This  displacement  is  entirely  independent  of  the  character 
of  the  salt,  provided  it  is  some  simple  one,  the  reaction  being  purely 
idnic  in  character.  The  place  of  hydrogen  in  the  series  is  most  in- 
teresting. All  those  metals  which  precede  it  will,  under 


NEUTRALIZATION  159 


conditions,  evolve   hydrogen   from    dilute   ar.ifls,  while   thosp. 
follow~wlll  not. 

From  the  standpoint  of  modern  electrical  theory  this  list  really 
represents  the  relative  ease  with  which  the  various  atoms  give  up 
one  or  more  electrons  to  form  ions.  Caesium,  the  metal  going  into 
solution  most  readily,  parts  with  an  electron  most  easily,  while  such 
metals  as  gold  and  platinum  retain  their  normal  number  of  electrons 
most  tenaciously. 

Nomenclature  of  acids,  bases,  and  salts.  The  naming  of  acids,  bases, 
and  salts  has  been  fairly  well  systematized,  although  there  are  some 
inconsistencies,  and  a  few  old  and  confusing  names  are  still  in  use. 

Acids.  As  a  rule,  the  acids  are  named  with  reference  to  their  oxygen 
content.  Those  which  contain  no  oxygen  are  all  called  Ay^ro-acids, 
the  name  ending  in  the  suffix  -ic  ;  thus,  hydrochloric  (HC1),  hydrosul- 
furic  (H2S),  hydrocyanic  (HNC).  Among  the  oxygen  acids  of  a  given 
element  some  one  is  chosen  and  given  a  name  suggestive  of  the  ele- 
ment and  ending  in  -ic,  but  with  no  prefix  ;  thus,  nitric  acid  (HNO3), 
sulfuric  (H2SO4),  chloric  (HC1O3).  This  name  is  usually  given  to  the 
most  common  acid  derived  from  the  element  in  question,  but  with 
rarer  elements  analogy  with  better-known  acids  is  usually  the  guide. 
The  other  acids  are  named  with  reference  to  this  one  as  standard.  One 
with  more  oxygen  is  given  a  name  with  the  prefix  per-  ;  thus,  perchloric 
acid  (HC1O4).  One  with  less  oxygen  is  named  with  the  suffix  -cms,  as 
nitrous  acid  (HNO2),  chlorous  (HC1O2),  sulfurous  (H2SO3).  One  with 
still  less  oxygen  is  given  the  prefix  hypo-,  as  hypochlorous  (HC1O). 

Bases.  The  bases  are  in  general  called  hydroxides,  the  name  of  the 
metal  being  prefixed  ;  thus,  potassium  hydroxide  (KOH),  copper  hy- 
droxide (Cu(OH)2).  When  the  same  metal  forms  two  hydroxides,  the 
name  of  the  one  with  the  smaller  number  of  hydroxyl  groups  ends 
in  -ous,  while  the  other  name  ends  in  -ic;  thus,  cuprous  hydroxide 
(CuOH),  cupric  hydroxide  (Cu(OH)2). 

Salts.  The  salts  are  named  with  reference  to  the  acids  from  which 
they  are  derived.  The  salts  of  hydro-acids  have  names  ending  in  -ide  ; 
thus,  sodium  chloride  (NaCl),  potassium  cyanide  (KNC).  The  same 
ending  -ide  is  also  given  to  the  name  of  any  compound  consisting  of 
but'  two  elements  (called  binary  compounds),  irrespective  of  their 
character.  Thus,  we  have  magnesium  nitride  (Mg3N2),  sulfur  oxide 
(SO2).  Salts  derived  from  oxygen  acids  whose  names  end  in  -ic 
have  names  ending  in  -ate  ;  thus,  sodium  nitrate  (XaNO3),  potassium 


160 


GENERAL  CHEMISTRY 


sulfate  (K2SO4).  In  a  similar  way  the  suffix  -ous  changes  to  -ite  when 
we  pass  from  the  acid  to  the  salt;  thus,  calcium  sulfite  (CaSO3), 
potassium  hypochlorite  (KC1O).  With  a  little  experience  these  prin- 
ciples will  become  familiar.  The  following  table  will  still  further 
illustrate  the  general  method  of  naming  acids  and  salts.  It  gives  the 
names  and  formulas  of  the  oxygen  acids  of  chlorine,  as  well  as  of 
the  sodium  salts  derived  from  each  of  these  acids. 


ACIDS 

SALTS 

Hypo-chlor-ous 

.     HC1O 

Sodium,  hypo-chlor-ite 

NaCIO 

Chlor-ous 

.     HC1O 

Sodium  chlor-it6 

NaCIO 

Chlor-ic  

.     HC1O3 

Sodium  chlor-ate 

.     NaClO3 

Per-chlor-ic  

.     HC1O4 

Sodium  per-chlor-ate 

.     NaClO4 

CHAPTER  XIV 

VALENCE  AND  STRUCTURAL  FORMULAS 

Valence  defined.  An  examination  of  the  formulas  which  have  been 
calculated  for  the  various  bases  brings  into  view  a  very  striking 
property  of  the  metallic  elements.  In  the  series 

K(OH)  Ca(OH)2  Al(OH),  Sn(OH)4 

it  is  at  once  seen  that,  while  the  atom  of  potassium  is  combined  with 
one  hydroxyl  group,  an  atom  of  calcium  is  combined  with  two,  an 
atom  of  aluminium  with  three,  and  one  of  tin  with  four.  A  series  of 
the  compounds  of  hydrogen  with  the  nonmetals  shows  the  same 
peculiarity : 

C1H  OH2  NH8  CH4 

With  oxides  there  is  an  even  greater  variety  in  numerical  ratios : 
K20      CaO      A1203      Sn02      P2O_      SO3      Mn2O7      OsO4 

1  2  3'4-5  6  7  8 

The  atom  of  each  element  evidently  possesses  some  property  which 
determines  how  many  atoms  of  another  kind  it  can  hold  in  combination. 
This  property  is  called  the  valence  of  an  atom  or  an  element. 

It  will  be  recalled  that  -the  equivalent  weight  of  an  element  is  the  weight 
which  combines  with  8  g.  of  oxygen,  and  that  either  the  equivalent  itself  or 
some  multiple  of  it  represents  the  relative  atomic  weight.  If  the  elements  were 
all  alike  in  the  number  of  their  atoms  which  combine  with  those  of  other  ele- 
ments, they  would  all  have  the  same  valence,  and  the  equivalent  weights  would 
then  represent  the  relative  atomic  weights  without  any  multiplication.  It  is  this 
peculiarity  of  valence  which  makes  it  necessary  to  multiply  some  equivalents  by 
an  integer  to  get  the  atomic  weight.  .Jjom  this  point  of  view  the  valence  of  an 
element  may  be  defined  as  the  number  of  equivalent  weights  of  the  element 
contained  in  its  atomic  weight.  

Unit  of  valence.  In  adopting  some  standard  unit  for  the  designation 
of  valence  it  must  be  clearly  understood  that  we  are  not  dealing  with 
the  intensity  of  attraction  between  the  elements.  Judging  by  the  heat  of 
combination,  hydrogen  has  a  far  stronger  affinity  for  oxygen  than  for 
nitrogen,  yet  an  atom  of  oxygen  can  unite  with  but  two  hydrogen  atoms, 

161 


^  ___  ^ 
/ 


162  GENERAL   CHEMISTRY 

while  one  of  nitrogen  can  combine  with  three.  Valence  has  to 

the  number  of  atoms  a  given  atom  can  hold  in  combination. 

As  a  standard  unit  some  element  should  therefore  be  selected  whose 
atom  never  holds  more  than  one  of  any  other  kind.  Such  an  element 
is  hydrogen,  and  the  hydrogen  atom  is  said  to  be  univalent.  An  atom 
or  radical  which  in  turn  combines  with  but  one  hydrogen  atom  is  like- 
wise univalent.  Examples  of  such  elements  are  chlorine,  iodine, 
potassium,  silver,  and  the  radicals  NH4  and  OH.  Similarly,  an  atom 
(or  radical)  that  holds  in  combination  two  univalent  -atoms,  as  oxygen 
in  the  compound  H2O,  is  said  to  be  divalent.  Likewise  we  have  triva- 
lent  elements,  as  nitrogen  in  the  compound  NH3,  and  tetravalent  ele- 
ments, as  carbon  in  the  compound  CH4.  There  seems  to  be  no  case 
of  valence  higher  than  8,  and  very  few  cases  where  it  exceeds  7. 

Referring  to  the  list  of  oxides  whose  formulas  are  given  on  page  161, 
it  will  be  seen  that  potassium  is  univalent  in  the  compound  K2O,  for 
the  2  atoms  of  potassium  must  have  a  total  valence  equal  to  that 
of  the  oxygen  atom  with  which  they  are  combined,  which  is  2.  One 
atom  of  potassium,  therefore,  has  a  valence  of  1.  Similarly,  calcium 
in  the  compound  CaO  is  divalent.  In  the  case  of  aluminium  oxide 
(A12O3)  the  2  atoms  of  aluminium  have  a  total  valence  of  6,  since 
they  hold  in  combination  3  divalent  oxygen  atoms.  One  atom  of 
aluminium,  therefore,  has  a  valence  of  3.  This  is  in  accord  with 
the  formula  for  the  chloride  A1C13.  By  the  same  method  of  reason- 
ing, each  of  the  elements  combined  with  oxygen  in  the  series  of 
oxides  just  mentioned  will  be  seen  to  have  the  valence  indicated  by 
the  numeral  placed  below  its  formula. 

Variable  valence.  At  one  time  it  was  thought  that  the  valence  of 
an  element  or  radical  is  always  the  same,  but  it  is  now  known  that 
nearly  all  of  them  may  have  different  valences. 

1.  Variation  of  valence  toward  the  same  element.  The  law  of  multiple 
proportion  reminds  us  that  one  element  may  combine  with  another  in 
everal  ratios,  which  is  merely  another  way  of  stating  the  fact  that  it 
may  have  more  than  one  valence.  In  the  oxides  CO  and  CO2  carbon 
acts  as  a  divalent  and  tetravalent  element  respectively.  In  the  oxides 
SO0  and  SOQ  sulfur  acts  as  a  tetravalent  and  hexavalent  element. 

£  o 

Frequently  a  metal  will  form  two  or  more  hydroxides,  as  Fe(OH)2 
and  Fe(OH)3,  each  giving  rise  to  a  series  of  salts.  In  such  cases  it  is 
usually  true  that  one  series  is  decidedly  more  stable  than  the  other, 
showing  that  the  valence  is  normally  that  of  the  more  stable  series. 


VALENCE  AND   STRUCTURAL  FORMULAS  163 

2.  Variations  of  valence  toward  different  elements.  If  we  regard  hydro- 
gen and  chlorine  as  the  standard  of  valence,  and  determine  the  valences 
of  the  elements  from  then*  compounds  with  these,  we  shall  in  general 
reach  different  conclusions  than  by  reasoning  from  the  formulas  of  the 
oxides.  The  following  examples  will  illustrate  this : 

C1H          SH2         NH3         CH4 
010.         S(X         NO.        (XX 


In  general  the  character  of  the  element  with  which  a  given  element 
is  combined  has  a  marked  influence  upon  its  valence.    An  element  is 
apt  to  show  its  lowest  valence  toward  hydrogen  or  chlorine  and  its 
highest  toward  oxygen.    Accordingly,  we  often  speak  of  the  oxygen  I     - 
valence  or  the  hydrogen  valence  of  an  element,  emphasizing  the  fact  J  "^ 
that  valence  is  a  relative  term. 

Representation  of  valence.  It  is  often  desirable  to  represent  in  some 
graphic  way  the  valence  of  an  element  present  in  a  compound, 
and  various  convenient  methods  have  been  employed.  The  usual 
method  is  to  draw  lines  between  the  symbols,  each  line  representing  a 
unit  of  valence.  Thus,  the  formula  for  hydrochloric  acid  is  sometimes 
written  H— Cl,  the  line  indicating  that  both  hydrogen  and  chlorine  are 
univalent  in  this  compound.  Likewise,  the  formula  for  water  may  be 

TT 

written   H— O— H  or  O<H,  indicating  that  oxygen  is  divalent  and 
hydrogen  univalent.   Again,  the  formula  for  aluminium  oxide  (Al2Og) 

Al^° 

is  sometimes  written  Tj>O  or  O=A1— O— A1=O,  indicating  that  in 

AHq 

this  compound  aluminium  is  trivalent  and  oxygen  divalent. 

This  method  of  indicating  valence  is  very  convenient,  but  it  may  lead  to 
inaccurate  language  and  erroneous  ideas.  Thus,  we  sometimes  say  that  two  ele- 
ments are  joined  or  united  by  two  valences,  as  in  the  case  of  calcium  oxide. 
This  is  a  brief  but  not  very  accurate  way  of  stating  the  fact  that  both  calcium 
and  oxygen  are  known  to  be  divalent,  and  that  the  oxide  is  represented  by  the 
formula  Ca  =  O.  It  is  not  valence  which  unites  these  elements,  but  chemical 
affinity. 

The  use  of  lines  to  represent  valence  is  also  open  to  the  objection  of  convey- 
ing the  impression  that  the  various  atoms  are  rigidly  united  in  some  way,  and 
this  impression  is  strengthened  by  the  fact  that  these  lines  are  usually  called 
bonds,  or  links.  There  is  no  evidence  for  any  such  rigid  union,  nor  can  we  well 
imagine  how  there  can  be  any.  On  the  other  hand,  we  have  every  reason  for 
thinking  that  the  atoms  in  a  molecule  maintain  an  orderly  relation  to  each  other, 
much  as  do  the  members  of  a  solar  system,  though  they  are  in  no  way  directly 
connected. 


164  GENERAL  CHEMISTRY 

Structural  formulas.  Formulas  such  as  those  just  mentioned,  with 
lines  indicating  the  valence  of  each  element,  are  called  structural,  or 
graphic,  formulas.  In  these  formulas  an  attempt  is  made,  in  a  general 
way,  to  represent  the  structural  relation  of  the  atoms  in  the  molecule. 
The  determination  of  the  arrangement  or  relative  grouping  of  the  atoms 
within  the  molecule  of  any  compound  is  based  upon  a  study  of  the 
general  properties  and  chemical  conduct  of  the  compound.  The  prob- 
lem is  not  a  simple  one,  and  its  solution  depends  upon  many  assump- 
tions, so  that  the  structural  formulas  of  many  compounds  are  unknown, 
while  in  other  cases  the  formulas  assigned  are  very  uncertain.  Never- 
theless, it  will  be  found  that  they  are  often  of  great  value. 

Determination  of  valence.  The  valence  of  an  element  may  usually 
be  inferred  by  noting  the  formula  of  its  compound  with  some  element 
of  known  valence,  such  as  hydrogen,  chlorine,  oxygen,  or  with  the 
hydroxyl  group.  For  example,  if  an  element  R  is  found  to  have  an 
oxide  R2O3,  a  hydroxide  R(OH)3,  a  chloride  RC13,  or  a  hydride  RH3, 
it  may  at  once  be  inferred  that  the  element  is  trivalent. 

The  valence  may  also  be  inferred  by  noting  the  number  of  hydro- 
gen atoms  which  one  atom  of  the  element  displaces  in  a  compound. 
For  example,  if  it  is  found  that  sodium  gives  the  sulfate  Na2SO4, 
the  sodium  is  univalent.  In  like  manner  the  formula  CaSO4  shows 
calcium  to  be  divalent,  while  the  formulas  A12(SO4)3  and  Sn(SO4)2 
show  aluminium  and  tin  to  be  trivalent  and  tetravalent  respectively. 

In  many  cases,  however,  more  information  is  needed  than  is  sup- 
plied by  the  formula.  For  example,  we  cannot  tell  the  valence  of 
sulfur  in  hydrogen  sulfate  (H2SO4)  until  we  know  how  the  sulfur 
atom  is  related  to  the  other  atoms  in  the  molecule.  Regarding  hydro- 
gen as  univalent  and  oxygen  as  divalent,  the  atoms  composing  the 
molecule  of  hydrogen  sulfate  may  have  various  arrangements,  as  sug- 
gested by  the  formulas  which  follow,  the  valence  of  the  sulfur  in 
each  of  these  being  indicated  by  the  number  placed  under  the  formula  : 


H-0     s  H-0-0     s 

H-0>    <*      H_0-0> 


10 


The  chemical  conduct  of  the  compound  indicates  that  the  structural 
formula  -j  Zo>sto  correctty  represents  the  relation  of  the  seven 
atoms  constituting  the  molecule,  so  that  the  sulfur  atom  in  hydrogen 
sulfate  is  hexavalent. 


VALENCE  AND   STRUCTURAL  FORMULAS  165 

Even  in  compounds  containing  but  two  elements  it  is  not  always 
possible  to  draw  correct  conclusions  from  the  mere  formula.  In  the 
compound  represented  by  the  formula  Fe3O4,  known  as  magnetic 
oxide  of  iron,  if  we  assume  oxygen  to  have  its  usual  valence  of  2, 
the  4  atoms  give  a  total  of  8.  This  is  distributed  among  3  atoms  of 
iron,  which  gives  a  fractional  value  of  2f  for  each.  Experiment  shows 
that  the  magnetic  oxide  is  really  a  union  of  two  distinct  oxides  of  the 
formulas  Fe  O  and  FeO,  in  the  first  of  which  the  iron  is  trivalent,  and 

2      3 

in  the  second,  divalent.  To  represent  this  fact  the  formula  is  often 
written  Fe2O3  •  FeO.  There  are  experimental  grounds  for  thinking  that 
the  structure  of  the  compound  is  correctly  represented  by  the  formula 

0  =  Fe-0       F 
O  =  Fe  -  O  ^ 

In  like  manner  we  were  unable  to  decide  as  to  the  valence  of  oxygen 
in  hydrogen  peroxide  until.it  was  shown  that  the  structure  of  the  com- 
pound is  represented  by  the  formula  H-O-Q-H,  in  which  the  valence 
of  oxygen  is  seen  to  be  2. 

Valence  of  radicals.  The  radical  SO4  is  evidently  divalent,  since 
it  combines  with  2  hydrogen  atoms,  or  with  1  of  calcium.  In  like 
manner  the  radicals  which  form  the  anions  of  the  various  acids  may 
be  considered  as  having  a  valence  equal  to  the  number  of  negative 
charges  which  they  carry,  or  to  the  number  of  hydrogen  ions  with 
which  they  unite.  Thus,  NO3  (derived  from  HNO3)  is  univalent ;  PO4 
(from  H3PO4)  is  trivalent ;  SiO4  (from  H4SiO4)  is  tetravalent.  The 
hydroxyl  radical  is  univalent,  as  shown  by  the  formula  K(OH). 

Molecular  formulas.  The  formula  Fe2Og-FeO  is  sometimes  called 
a  molecular  formula,  since  it  represents  the  union  of  two  molecules, 
each  of  which  is  known,  in  the  free  state,  as  an  independent  sub- 
stance. Salts,  containing  water  of  crystallization,  such  as  the  hydrate 
CuSO4-5H2O,  are  given  similar  formulas.  In  many  such  cases  we 
have  no  knowledge  as  to  the  structure  of  the  compound,  and  the 
formula  is  intended  to  be  entirely  noncommittal.  It  does  not  mean 
that  the  two  kinds  of  molecules  unite  without  any  rearrangement, 
but  merely  that  at  present  we  have  no  information  as  to  its  character. 
As  our  knowledge  increases,  such  formulas  gradually  give  way  to  real 
structural  ones. 

Nature  of  valence.  We  have  very  little  knowledge  as  to  why  the 
potassium  atom  can  hold  but  one  chlorine  atom,  while  that  of  calcium 
can  hold  two  and  that  of  aluminium  three.  So  long  as  we  know  little 


166  GENERAL  CHEMISTRY 

about  the  nature  of  chemical  affinity,  which  causes  the  union,  we  cannot 
hope  to  have  clear  ideas  about  valence,  which  is  merely  the  numerical 
ratio  in  which  atoms  combine.  It  seems  probable  that  as  we  learn 
more  about  electricity,  both  chemical  affinity  and  valence  will  be 
much  better  understood,  since  all  three  appear  to  be  very  intimately 
related  to  each  other. 

Applications  of  valence.  A  knowledge  of  the  valence  of  an  element 
is  of  great  assistance  in  recalling  the  formulas  of  its  compounds.  The 
fact  that  calcium  is  divalent  as  measured  by  the  hydrogen  it  displaces 
from  acids  at  once  enables  us  to  write  the  probable  formulas  of  all  the 
salts  of  calcium,  provided  the  formulas  of  the  acids  are  known.  The 
following  formulas  will  make  this  clear : 

HC1  H2(S04)  H3(P04)  H4(Si04) 

CaCl2  CaSO4  Ca3(PO4)2  CaaSiO4 

In  writing  equations,  the  same  principles  must  be  observed.  For 
example,  in  the  general  reaction  of  an  acid  upon  a  base  to  form  a  salt, 
the  essential  action  is  the  combination  of  the  hydroxyl  with  the  hydro- 
gen ions.  The  acid  and  base  must  therefore  be  taken  in  such  pro- 
portions as  will  give  the  same  number  of  each  of  these  ions.  Thus, 
in  the  reaction  of  ferric  hydroxide  (Fe(OH)3)  upon  sulfuric  acid 
(H2SO4)  it  will  be  necessary  to  take  2  molecules  of  Fe(OH)3  and 
3  of  H2SO4: 

2  Fe(OH)3  +  3  H2SO4  =  Fea(SO4)8  +  6  H2O 


CHAPTER  XY 

COMPOUNDS  OF  NITROGEN 

Occurrence.  Although  large  quantities  of  nitrogen  occur  in  the  at= 
mosphere,  it  is  all  in  the  free  state,  with  the  exception  of  a  com- 
paratively small  amount,  which  is  present  in  the  form  of  ammonia 
and  oxides  of  nitrogen,  or  compounds  derived  from  these.  In  the 
materials  composing  the  earth's  crust,  on  the  other  hand,  there  oc- 
cur in  certain  localities  considerable  deposits  of  compounds  of  nitro- 
gen, especially  of  sodium  nitrate  (XaNO3)  and  potassium  nitrate 
(KNOg).  Moreover,  such  compounds  are  present,  at  least  in  small 
quantities,  in  all  productive  soils.  From  these  soils  the  nitrogen  is 
taken  up  by  plants  and  built  into  complex  compounds.  Animals  feed- 
ing on  these  plants  assimilate  the  nitrogenous  matter,  which  thereby 
becomes  an  essential  part  of  the  animal  tissue.  In  both  plants  and 
animals  the  nitrogen  is  present  chiefly  in  the  form  of  protein  mat- 
ter, which  consists  of  complex  compounds  containing  the  elements 
carbon,  nitrogen,  oxygen,  and  hydrogen,  and  sometimes  phosphorus 
and  sulfur. 

The  unstable  character  of  compounds  of  nitrogen.  Experiment  shows 
that  the  molecule  of  nitrogen  has  the  formula  X2,  and  that  the  ele- 
ment is  very  inactive  at  ordinary  temperatures.  This  inactivity  seems 
to  be  partly  due  to  the  fact  that  the  nitrogen  molecule  is  quite  stable 
and  that  a  good  deal  of  energy  is  required  to  separate  it  into  its  atoms, 
into  which  form,  as  a  rule,  it  is  converted  before  entering  into  com- 
bination with  other  elements.  On  the  other  hand,  when  nitrogen 
occurs  as  a  constituent  of  a  compound,  the  nitrogen  atoms  tend  to 
leave  the  compound  and  form  stable  nitrogen  molecules.  As  a  result 
of  this  tendency  compounds  containing  nitrogen  are  apt  to  be  un- 
stable. It  is  partly  due  to  the  unstable  character  of  certain  nitrog- 
enous compounds  that  they  are  so  extensively  used  as  a  constituent 
of  explosives. 

While  a  great  many  compounds  of  nitrogen  are  known,  it  is  desira- 
ble at  this  time  to  discuss  only  some  of  the  more  simple  ones,  namely, 
those  which  nitrogen  forms  with  hydrogen  and  oxygen. 

167 


168  GENERAL  CHEMISTRY 

COMPOUNDS  OF  NITROGEN  WITH  HYDROGEN 

Nitrogen  forms  three  important  compounds  with  hydrogen,  the 
names  and  formulas  of  which  are  as  follows  :  ammonia  (NHg),  a  color- 
less, gaseous  compound  of  characteristic  odor;  hydrazine  (N2H4),  a 
colorless  liquid  boiling  at  113°;  hydronitric  acid  (HN3),  a  very  un- 
stable, colorless  liquid  boiling  at  37°.  These  compounds  will  now 
be  discussed  more  or  less  in  detail,  according  to  their  importance. 

Ammonia.  Inasmuch  as  ammonia  is  formed  in  certain  natural 
processes  which  are  constantly  taking  place  about  us,  such  as  the 
decay  of  nitrogenous  organic  matter,  it  is  easy  to  understand  why 
this  compound  has  been  known  for  so  long  a  time.  It  was  originally 
prepared  by  heating  such  tissues  as  the  hoofs  and  horns  of  animals, 
and  the  aqueous  solution  of  the  gas  so  obtained  was  termed  spirits 
of  hartshorn.  The  pure  gas  itself  was  first  prepared  by  Priestley,  in 
1774,  and  its  composition  was  determined  soon  after  by  the  French 
chemist  Berthollet. 

Preparation  of  ammonia.  Ammonia  can  be  prepared  in  a  number  of 
ways,  the  most  important  of  which  are  the  following  : 

1.  Synthetic  method.  When  electric  sparks  are  passed  through  a 
mixture  of  nitrogen  and  hydrogen,  small  amounts  of  ammonia  are 
formed  by  the  direct  union  of  the  elements.  The  limited  yield  is  due 
to  the  fact  that  the  reaction  is  reversible,  as  indicated  below,  and  equi- 
librium results  when  a  relatively  small  amount  of  ammonia  has  been 

formed: 


2.  Laboratory  method.  In  the  laboratory,  ammonia  is  usually  pre- 
pared from  ammonium  chloride  (NH4C1),  a  white  solid  obtained  in 
the  manufacture  of  coal  gas.  In  this  compound  the  group  of  atoms 
NH4  acts  as  a  univalent  radical  and  is  termed  ammonium  ;  hence  the 
name  ammonium  chloride.  When  a  mixture  of  ammonium  chloride  and 
sodium  hydroxide  is  heated  in  the  presence  of  a  small  quantity  of 
water,  the  ammonium  radical  and  sodium  change  places,  as  expressed 
in  the  following  equation  : 

NH4C1  +  NaOH  =  NaCl  +  NH4OH 

The  resulting  ammonium  hydroxide  (NH4OH)  is  unstable  and,  as  fast 
as  it  is  formed,  breaks  down  into  water  and  ammonia  : 

NH  OH  =  NH.  +  HO 


COMPOUNDS  OF  NITROGEN 


169 


Calcium  hydroxide  (Ca(OH)2)  is  frequently  used  in  place  of  the 
more  expensive  sodium  hydroxide  : 


2  NH4C1+  Ca(OH)2  =  CaCl 


2  NH/)H 


IK) 


FIG.  68 


The  preparation  is  conducted 
as  follows  :  The  mixture  of  am- 
monium chloride  and  calcium  hy- 
droxide is  introduced  into  a  flask 
with  a  little  water  (Fig.  68,  A) 
and  gently  heated.  Ammonia  is 
evolved  and  may  be  collected  by 
bringing  the  end  of  the  exit  tube 
inside  and  near  the  bottom  of  an 
inverted  bottle  B,  as  shown  in 
the  figure.  The  gas,  being  lighter 
than  air,  collects  in  the  bottle, 
gradually  forcing  the  air  out  at 
the  mouth.  Because  of  its  great 
solubility  it  is  not  practicable  to 
collect  ammonia  over  water,  as 
in  the  case  of  oxygen  and  hydro- 
gen. In  place  of  water,  however, 

one  may  use  some  liquid,  such  as  mercury,  in  which  the  ammonia  is  not  soluble. 

As  ordinarily  prepared  in  the  laboratory,  it  is  collected  by  simply  displacing  the  air 

in  a  bottle,  as  described 

above. 

In  addition  to  this 
method,  ammonia  is 
often  obtained  in  the 
laboratory  by  heating 
an  aqueous  solution 
of  the  gas.  Such  a 
solution  is  a  common 
article  of  commerce, 
and  is  sold  under  the 
name  ammonia  water, 
or  aqua  ammonia.  The  - 
solution  is  placed  in 
a  flask  A  (Fig.  69)  and  a  gentle  heat  applied.  Ammonia  is  evolved 
and  is  passed  through  a  cylinder  B  filled  with  small  pieces  of  lime, 
which  serve  to  remove  any  moisture.  The  pure  gas  so  obtained  may 
be  collected  by  the  displacement  of  air,  as  described  above. 


FIG.  69 


170 


GENERAL  CHEMISTRY 


Commercial  preparation.  Ammonia  is  obtained  commercially  in  the 
process  of  manufacturing  coal  gas.  Certain  grades  of  bituminous,  or 
soft,  coal  are  best  adapted  for  this  purpose.  Such  coal  contains,  in  addi- 
tion to  carbon,  about  1  per  cent  of  nitrogen  and  7  per  cent  of  hydrogen, 
as  well  as  small  percentages  of  other  elements.  When  such  coal  is  heated 
in  retorts  from  which  the  air  is  excluded  (Fig.  112,  p.  323),  complicated 
changes  take  place,  resulting  not  only  in  the  formation  of  the  combusti- 
ble gases  which  constitute  coal  gas,  but  also  of  ammonia  and  many  other 
valuable  products.  From  25  to  50  per  cent  of  the  nitrogen  present 
in  the  coal  is  converted  into  ammonia.  The  volatile  matter  expelled 
from  the 'coal  is  passed  through  water,  which  absorbs  the  ammonia, 
together  with  certain  other  compounds,  forming  a  solution  known  as 
the  ammoniacal  liquor.  When  this  liquor  is  heated  with  slaked  lime, 
ammonia  is  evolved  and  may  be  passed  into  pure  cold  water,  forming 
ordinary  aqua  ammonia ;  or  it  may  be  passed  into  dilute  solutions  of 
either  hydrochloric  or  sulfuric  acid,  forming  ammonium  chloride  and 
ammonium  sulfate  respectively.  The  reactions  which  result  in  the 
formation  of  these  compounds  will  be  discussed  later  in  this  chapter. 

Properties.  Ammonia  is  a  colorless  gas  having  a  strong,  suffocating 
odor.  Under  standard  conditions  1 1.  of  the  pure  gas  weighs  0.7708  g., 
being  0.59  times  as  heavy  as  air.  Its  critical  temperature  is  131°,  at 
which  temperature  it  is  liquefied  by  a  pressure  of  113  atmospheres. 
Liquid  ammonia  is  colorless  and  boils  at  —  33.5°.  The  properties  of 
liquid  ammonia  have  been  extensively  studied  by  E.  C.  Franklin,  who 
has  shown  that,  like  water,  it  is  not  only  an  excellent  solvent  but  also 
a  highly  ionizing  one.  At  still  lower  temperatures  ammonia  can  be 
obtained  in  the  form  of  a  snowlike  solid  melting  at  —  75.5°. 

A  noteworthy  property  of  ammonia  is  its  extreme  solubility  in 
water.  Under  a  pressure  of  1  atmosphere,  1  1.  of  water  dissolves 
1298  1.  of  the  gas  at  0°,  and  710  1.  at  20°.  On  account  of  the  expan- 
sion of  the  liquid  the  resulting  solutions  have  a  density  less  than  that 
of  water,  as  is  shown  in  the  following  table : 


Density  .... 

1.00 

0.995 

0.990 

0.980 

0.970 

0.950 

0.930 

0.910 

0.890 

0.880 

Per  cent  by  weight 

ofNH8    .     .     . 

0.00 

1.14 

2.31 

4.80 

7.31 

12.74 

18.64 

24.99 

31.73 

35.60 

Chemical  conduct.   The   chemical  properties   of  ammonia  may   be 
conveniently  discussed  under  the  following  heads: 


\ 


COMPOUNDS  OF  NITROGEN 


1T1 


1.  Dissociation  of  ammonia.    At  ordinary  temperatures  ammonia  is 
a  stable  compound.    When  heated  to  a  high  temperature,  however, 
or  when  subjected  to  the  action  of  electric  sparks,  it  is  dissociated 
into  its  elements.     The  reaction  is  a  reversible  one.     Thus,  when 
electric  sparks  are  passed  through  ammonia,  equilibrium  results  when 
about  96  per  cent  of  the  compound  is  dissociated  (see  synthetic  method 
of  preparation,  p.  168). 

2.  Reducing  action  of  ammonia.    When  heated,  ammonia  acts  as  a 
strong  reducing  agent,  owing  to  the  hydrogen  which  is  liberated  from 
the  compound  under  the  influence  of  heat.   This  reducing  property 
may  be  shown  by  passing  the  gas  through  a  tube  containing  a  metallic 
oxide,  such  as  oxide  of  copper,  heated  to 

a  high  temperature.   The  oxide  is  reduced 
to  the  metal : 

2  NH3  +  3  CuO  =  3  Cu  +  N2  +  3  H2O 

3.  Relation  to  combustion.  Because  of  the 
hydrogen  present  in  ammonia,  and  because 
of  the  comparative  instability  of  the  com- 
pound, the  gas  readily  burns  in  an  atmos- 
phere of  pure   oxygen.    The  combustion 
will  not  take  place  in  air,  however,  unless 
heat  is  continuously  applied  from  an  ex- 
ternal source. 

These  facts  may  be  shown  in  the  following 
way:  Some  aqua  ammonia  is  poured  into  the 
flask  A  (Fig.  70)  and  heated  gently.  Ammonia 
escapes  and  may  be  ignited  by  holding  a  Bunsen 
flame  at  the  end  of  the  exit  tube  B.  Combustion 

ceases,  however,  as  soon  as  the  flame  is  withdrawn.  If  now  oxygen  is  passed  in 
through  the  tube  C,  the  ammonia  escaping  from  B  is  surrounded  by  pure  oxygen 
and,  if  ignited,  will  continue  to  burn. 

4.  Action  upon  metals.    A  number  of  the  metals,  such  as  magnesium 
and  lithium,  react  with  ammonia  at  high  temperatures,  displacing  all 
the  hydrogen  and  forming  nitrides : 

3  Mg  +  2  NH3  =  Mg3N2  +  3  H2 
These  nitrides  are  solids  and  react  with  water  to  form  ammonia : 


FIG.  70 


Mg3X2  +  6  H20  =  3  Mg  (OH),  +  2  NH, 


172  GENERAL  CHEMISTRY 

On  the  other  hand,  the  metals  sodium  and  potassium  act  upon  am- 
monia, displacing  only  a  part  of  the  hydrogen : 

2  Na  +  2  NH3  =  2  NaNH2  +  H2 

The  radical  NH2  is  termed  the  amido,  or  amino,  group,  so  that  the  com- 
pound NaNH2  is  termed  sodamide.  It  is  a  yellowish  solid  and,  like 
the  nitrides,  reacts  with  water  to  form  ammonia. 

5.  Action  upon  salts.    Ammonia  combines  directly  with  a  number 
of  salts  to  form  complex  compounds.   In  some  of  these  compounds, 
such  as  that  represented  by  the  formula  CaCl2 ••  8  NH3,  the  ammonia 
seems  to  play  much  the  same  part  as  does  the  water  in  hydrates. 

6.  Action  toward  water.  It  will  be  recalled  that  ammonia  is  extremely 
soluble  in  water,  and  that  the  resulting  solution  (aqua  ammonia)  is 
basic  in  character  (p.  125).  These  properties  are  accounted  for  by 
the  fact  that  ammonia  and  water  enter  into  chemical  combination  with 
each    other    and   that   the   resulting   compound,  namely,  ammonium 
hydroxide  (NH4OH),  ionizes   (p.  154),  forming  the  two  univalent 
ions  NH4+  and  OH~,  the  latter  imparting  to  the  solution  its  basic 
character.    In  this  solution  the  ammonium  hydroxide  is  in  equilibrium 
not  only  with  ammonia  and  water  but  also  with  the  ions  NH4+  and 
OH~,  as  expressed  in  the  following  equation : 

NH3  +  H20  +±L  NH4OH  ^=±  NH4+,  OH~ 

When  heated,  aqua  ammonia  acts  as  if  it  were  solely  a  solution 
of  ammonia  in  water,  for  as  fast  as  the  ammonia  is  driven  out  by 
the  heat,  the  equilibrium  existing  among  the  different  substances 
represented  in  the  equation  is  disturbed,  with  the  result  that  more 
ammonia  is  formed.  If  the  heating  is  continued,  therefore,  all  of  the 
ammonia  is  finally  driven  out.  On  the  other  hand,  when  acted  upon 
by  an  acid,  aqua  ammonia  acts  as  if  it  were  solely  a  solution  of  am- 
monium hydroxide,  for  as  fast  as  the  ammonium  hydroxide  present  is 
neutralized  by  the  acid,  more  is  formed  as  long  as  any  free  ammonia 
remains  in  solution.  Thus,  when  hydrochloric  acid  is  added  to  aqua 
ammonia,  ammonium  chloride  and  water  are  formed : 

NH4+,  OH-  +  H+,  Cl-  =  NH4+,  Cl~  +  H2O 

On  evaporating  the  water,  the  ammonium  chloride  is  left  in  the  form 
of  a  white  solid.  Similarly,  nitric  and  sulfuric  acids  form  respectively 
ammonium  nitrate  (NH4NOg)  and  ammonium  sulfate  ((NH4)2SO4). 


COMPOUNDS  OF  NITROGEN  173 

These  compounds  are  also  formed  by  the  direct  combination  of 
ammonia  with  the  corresponding  acid: 

NH3  +  HC1  =  NH4C1 

NH3  +  HNO3  =  NH4NO8 

2NH3  +  H2S04  =  (NH4)2S04 

Composition  of  ammonia.  That  ammonia  is  a  compound  of  nitrogen 
and  hydrogen  is  proved  by  the  fact  that  it  may  be  formed  by  the 
direct  union  of  the  two  elements  (see  method  of  preparation).  The 
quantitative  composition  of  the  compound  may  be  determined  by  tak- 
ing advantage  of  certain  reactions  which  make  it  possible  to  liberate 
the  nitrogen  as  well  as  the  hydrogen  from  any  definite  volume  of  am- 
monia. By  measuring  the  volumes  of  the  gases  so  liberated  one  can 
compare  them  not  only  with  each  other  but  also  with  the  volume  of 
the  ammonia  from  which  they  were  derived.  In  this  way  it  has  been 
proved  that  2  volumes  of  ammonia  yield  on  decomposition  1  volume 
of  nitrogen  and  3  volumes  of  hydrogen,  as  expressed  graphically  in 
the  following  equation : 


2  vol.  of  NH3  =      1  vol.  of  N2        +        3  vol.  of  H2. 

Of  course,  the  reverse  is  likewise  true  :  3  volumes  of  hydrogen  com- 
bine with  1  volume  of  nitrogen  to  form  2  volumes  of  ammonia 
(compare  with  the  composition  of  water,  p.  69).  By  comparing  the 
weight  of  the  3  volumes  of  hydrogen  with  that  of  the  1  volume  of 
nitrogen  we  can  determine  the  proportion  by  weight  in  which  these 
two  elements  combine  to  form  ammonia.  The  results  are  expressed  by 
the  formula  NHg,  which  indicates  that  14.01  parts  by  weight  of  nitro- 
gen combine  with  3.024  of  hydrogen  to  form  17.034  parts  of  ammonia. 
Each  of  the  hydrogen  atoms  in  ammonia  apparently  bears  the  same 
relation  to  the  nitrogen  atom  as  is  represented  in  the  structural 
formula  I,  below.  The  nitrogen  is  therefore  trivalent  in  this  com- 
pound. When  ammonia  combines  with  an  acid,  such  as  hydrochloric 
acid,  to  form  salts,  the  nitrogen  becomes  pentavalent,  as  expressed  in 
the  structural  formula  II  for  ammonium  chloride. 

H  TT  .  H 


H 
I  II 


174 


GENERAL   CHEMISTRY 


Uses  of  ammonia.  Large  quantities  of  ammonia  are  used  in  the 
manufacture  of  aqua  ammonia,  as  well  as  in  the  formation  of  am- 
monium compounds,  such  as  ammonium  chloride  and  ammonium  sul- 
fate.  In  the  liquid  state  it  is  also  used  extensively  in  the  manufacture 
of  artificial  ice.  Its  use  for  this  purpose  is  based  on  the  facts  that  the 
gas  is  easily  liquefied  by  pressure  and  that  the  resulting  liquid  has  a 
relatively  high  heat  of  vaporization. 

The  manufacture  of  artificial  ice.  The  general  method  used  in  the  manufacture 
of  artificial  ice  may  be  understood  by  reference  to  Fig.  71.  Ammonia  is  liquefied 
by  means  of  a  compressor  pump  and  led  into  the  pipes  A,  B.  The  heat  of  con- 


1 


(H: 


FIG.  71 


densation  is  absorbed  by  water 
flowing  over  the  pipes.  These 
pipes  lead  into  coils  in  a  large 
tank  nearly  filled  with  brine, 
prepared  by  dissolving  calcium 
chloride  in  water.  By  means  of 
an  expansion  valve  C  the  pres- 
sure upon  the  liquid  ammonia 
is  diminished  as  it  enters  the 
coils,  and  the  heat  absorbed  by 
the  rapid  evaporation  of  the 
liquid  lowers  the  temperature 
of  the  brine  below  0°.  Metal 
vessels  D,  E,  F  filled  with  pure 
water  are  lowered  into  the  cold 
brine  and  left  until  the  water  in  them  is  frozen  into  cakes  of  ice.  The  gaseous 
ammonia  resulting  from  the  vaporization  is  led  through  G  back  to  the  com- 
pressor pump,  by  which  it  is  again  liquefied,  so  that  the  process  is  continuous. 

Hydrazine  (N2H4).  This  compound  may  be  regarded  as  ammonia 
in  which  an  atom  of  hydrogen  has  been  displaced  by  the  univalent 
radical  NH2.  Its  formula  may  therefore  be  written  NH2  —  NH2.  It  is 
formed  by  the  reduction  of  hyponitrous  acid  (H2N2O2),  but  is  most 
readily  obtained  by  complicated  reactions  with  certain  organic  com- 
pounds. When  an  aqueous  solution  is  distilled,  the  hydrate  (N2H4  •  H2O) 
is  obtained.  The  free  hydrazine  is  prepared  from  this  hydrate  by  dis- 
tilling it  under  diminished  pressure  with  some  compound,  such  as 
barium  oxide,  that  has  a  strong  affinity  for  water. 

Hydrazine  is  a  colorless  liquid  boiling  at  113.5°.  Like  ammonia,  it 
combines  with  water  to  form  a  base  from  which  salts  can  be  prepared 
by  the  action  of  acids.  For  example,  with  hydrochloric  acid  it  forms 
hydrazine  hydrochloride,  a  white  solid  having  the  formula  (N2H4  •  HC1). 
It  is  a  strong  reducing  agent. 


COMPOUNDS  OF  NITROGEN  175 

Hydronitric  acid  (HN3).  This  acid,  known  also  as  hydrazoic  acid, 
was  first  obtained  by  Curtius  in  1890,  and  the  properties  of  the  acid, 
as  well  as  its  salts,  have  been  extensively  studied  by  Dennis  and  his 
associates.  Its  sodium  salt  may  be  prepared  by  the  action  of  nitrous 
oxide  (a  gas  having  the  formula  N2O)  on  sodamide  : 


NaNH2  +  N2O  =  NaN3  +  H2O 

An  aqueous  solution  of  the  free  acid  may  be  obtained  from  the 
sodium  salt  by  adding  dilute  sulfuric  acid  and  distilling: 

NaN,  +  H2S04  =  HN3  +  NaHSO4 

The  pure  acid  is  a  colorless  liquid  of  disagreeable  odor.  It  boils  at 
37°  and  is  violently  explosive,  dissociating  into  its  constituent  ele- 
ments with  the  liberation  of  considerable  heat.  It  has  the  general 
properties  of  a  weak  acid,  dissolving  certain  metals,  such  as  zinc  and 
iron,  forming  the  corresponding  salts  and  evolving  hydrogen.  The 
salts  of  hydronitric  acid  are  solids,  most  of  them  being  colorless. 
Some  of  them  are  violently  explosive  and  all,  on  being  heated,  decom- 
pose, evolving  nitrogen  and  leaving  the  metal  uncombined.  The 
structural  formula  of  hydronitric  acid  is  not  known,  but  the  follow- 
ing has  been  suggested  :  ^ 

H-N<|| 
XN 

COMPOUNDS  OF  NITROGEN  WITH  HYDROGEN  AND  OXYGEN 

The  most  important  of  these  compounds  are  the  following: 
Hyponitrous  acid  (H2N2O2)  :  a  colorless,  unstable  solid. 
'Nitrous  acid  (HNO2)  :  known  only  in  dilute  solution. 
Nitric  acid  (HNO3)  :  a  colorless  liquid. 
Hydroxylamine  (NH2OH)  :  a  white,  crystalline  solid. 

The  first  three  of  these  compounds  are  acids,  as  is  indicated  by  the 
names  ;  the  last  is  a  base.  Nitric  acid  will  be  described  first,  since  it 
is  by  far  the  most  important  compound  of  the  group. 

Nitric  acid.  Nitric  acid  was  well  known  to  the  alchemists,  being  first 
prepared  by  the  Egyptians.  In  the  ninth  century  the  alchemist  Geber 
prepared  it  from  saltpeter  (KNO3)  by  a  process  somewhat  similar  to 
that  used  at  the  present  time,  and  the  Germans  still  call  it  Salpeter- 
saure.  The  composition  of  the  acid  was  first  determined  by  Lavoisier 
and  Priestley. 


176 


GENEBAL   CHEMISTRY 


Because  of  its  great  activity,  nitric  acid  does  not  occur  free  in 
nature,  but  a  number  of  its  salts  are  found  in  considerable  quanti- 
ties. The  most  abundant  of  these  is  sodium  nitrate,  which  is  found 
in  large  quantities  in  Chile,  and  hence  is  known  as  Chile  saltpeter. 

Preparation.  Nitric  acid  is  prepared,  both  in  the  laboratory  and 
commercially,  by  the  action  of  sulfuric  acid  upon  some  salt  of  nitric 
acid.  On  account  of  its  low  cost  the  salt  generally  used  is  sodium 
nitrate.  When  brought  together,  sulfuric  acid  and  sodium  nitrate 
react  to  form  sodium  hydrogen  sulfate  and  nitric  acid,  as  expressed 
in  the  following  equation  : 

NaNOa  +  HjSCX  ^zfc  NaHSO.  +  HNO 

o^4  4  3 

The  reaction  is  a  reversible  one,  and  at  ordinary  temperatures  equi- 
librium soon  results.  If,  however,  a  gentle  heat  is  applied  to  the 
mixture,  the  nitric  acid  is  removed  as  fast  as  formed,  since  it  has  a 
relatively  low  boiling  point  (86°),  and  may  be  recovered  by  con- 

densing the  vapors.  Un- 
der these  conditions  the 
reverse  reaction  cannot 
take  place,  and  the  ac- 
tion between  the  sodium 
nitrate  and  the  sulfuric 
acid  continues  until  it  is 
completed. 

If    double    the    weight  of 
sodium    nitrate   indicated  in 
FIG.  72  tne  equation  is  used,  and  the 

mixture  is  heated  to  a  higher 

temperature,  the  sodium  hydrogen  sulfate  which  is  first  formed  reacts  with  the 
excess  of  sodium  nitrate  to  form  the  normal  sodium  sulfate  and  nitric  acicfc 


NaIISO 


3  =  Na2SO4  +  HNO3 


This  additional  amount  of  nitric  acid  is  formed  without  using  any  additional  sul- 
furic acid.  The  higher  temperature  required  to  bring  about  the  reaction,  however, 
partly  decomposes  the  nitric  acid  so  that  the  process  is  not  an  economical  one. 

The  preparation  of  nitric  acid  in  the  laboratory  may  be  conducted  as  follows  : 
sodium  nitrate  is  introduced  into  a  retort  A  (Fig.  72)  and  sulfuric  acid  added. 
If  the  retort  is  gently  heated,  the  nitric  acid  distills  over  as  fast  as  formed, 
and  is  condensed  in  a  tube  B  kept  cool  by  ice  water. 

Commercial  preparation  of  nitric  acid.  Fig.  73  illustrates  a  form  of  apparatus 
used  in  the  preparation  of  nitric  acid  on  a  large  scale.  Sodium  nitrate  and  sul- 
furic acid  are  heated  in  the  iron  retort  A.  The  resulting  acid  vapors  pass  in 


COMPOUNDS  OF  NITROGEN 


177 


the  direction  indicated  by  the  arrows,  and  are  condensed  in  the  glass  tubes  B, 

which  are  covered  with  cloth  kept  cool  by  streams  of  water.    These  tubes  are  in- 

clined so  that  the  liquid  result- 

ing from  the  condensation  of  Water 

the  vapors  runs  back   into   C 

and  is  drawn  off  into  the  large 

vessel  Z). 

Because  of  the  unstable  char- 
acter of  nitric  acid  a  certain 
amount  of  decomposition  al- 
ways takes  place  when  it  is 
distilled  under  ordinary  atmos- 
pheric pressure.  To  prevent 
this  the  process  is  often  carried 
out  under  diminished  pressure. 

Formation  of  nitric  acid 
from  air.  When  electric 
sparks  are  passed  through  air,  a  portion  of  the  nitrogen  and  oxy- 
gen present  combine  to  form  a  colorless  gaseous  compound  known 
as  nitric  oxide,  which  has  the  formula  NO.  As  fast  as  it  is  formed, 
this  oxide  combines  with  more  oxygen  to  form  nitrogen  dioxide 
(NO^,  a  reddish-brown  gas.  When  the  dioxide  is  brought  in  con- 
tact with  water,  the  two  react  to  form  nitric  acid.  The  equations 
for  the  reactions  are  as  follows: 


Fro.  73 


22 

2  NO  +  O2  =  2  NO2 
3  N02  +  H20  =  2  HN08  +  NO 

From  the  acid  so  obtained,  its  various  salts  (the  nitrates)  may  be  pre- 
pared by  neutralization  with  the  appropriate  bases.  Inasmuch  as  large 
quantities  of  the  nitrates  are  in  demand  in  the  manufacture  of  fertil- 
izers, and  since  the  supply  of  these  compounds  in  nature  is  limited, 
repeated  efforts  have  been  made  to  prepare  both  nitric  acid  and  the 
nitrates  from  the  inexhaustible  supplies  of  oxygen  and  nitrogen  in  the 
atmosphere,  utilizing  the  above  reaction.  The  method  has  not  as  yet 
proved  a  marked  success,  on  account  of  the  cost  of  generating  the 
electric  current.  It  has  been  greatly  improved,  however,  by  Berkeland 
and  Eyde,  and  their  process  is  in  use  in  Norway,  where  the  waterfalls 
are  utilized  for  the  generation  of  electrical  energy. 

In  the  Berkeland  and  Eyde  process  an  electric  arc  is  produced  by  forcing  a 
powerful  alternating  current  between  two  copper  electrodes  so  placed  that  the 


GE1 


178  GENERAL  CHEMISTRY 

* 

arc  passes  between  the  poles  of  an  electromagnet   (Fig.  74).    By  the  action 

of  the  magnet  the  arc  is  spread  out  into  disks  about  2  m.  in  diameter.  The 

apparatus  is  inclosed  in  a  large  tube  in 
such  a  way  that  air  forced  through  the  tube 
is  subjected  to  the  action  of  the  electric 
discharge.  The  greater  efficiency  of  this 
process  lies  in  the  fact  that,  by  the  spread- 
ing out  of  the  arc  into  disks,  larger  quan- 
tities of  air  come  under  its  influence.  The 
nitrogen  dioxide  so  formed  is  passed  into 
water,  and  a  dilute  solution  of  nitric  acid 
is  obtained.  The  cost  attending  the  con- 
FIG.  74  centration  of  the  acid  is  so  great  that  it 

has  not  been  found  economical  to  prepare 

the  pure  acid  by  this  method,  but  the  dilute  solution  is  being  used  to  some 

extent  in  the  preparation  of  nitrates. 

Preparation  of  pure  nitric  acid  (hydrogen  nitrate).  Pure  nitric 
acid  (more  properly  called  hydrogen  nitrate,  to  distinguish  it  from  its 
aqueous  solutions)  readily  decomposes  into  water,  nitrogen  dioxide, 
and  oxygen,  as  represented  in  the  following  equation: 


4  HNO3  =  2  H2O 


4  NO  +  O 


The  nitrogen  dioxide  resulting  from  the  decomposition  is  a  reddish- 
brown  gas,  which  dissolves  in  the  liquid,  imparting  to  it  a  yellowish 
color.  Because  of  its  unstable  character,  hydrogen  nitrate  is  difficult 
to  prepare.  In  its  preparation  from  sulfuric  acid  and  sodium  nitrate  a 
slight  amount  of  decomposition  takes  place,  but  this  can  be  largely 
prevented  by  conducting  the  distillation  under  diminished  pressure, 
and  in  this  way  a  nearly  pure  compound  is  obtained.  If  this  solution 
is  cooled  to  a  low  temperature,  pure  hydrogen  nitrate  crystallizes  out 
in  the  form  of  a  snowlike  solid  melting  at  —  41.3°.  As  the  temperature 
rises,  the  solid  melts  to  a  colorless  liquid,  but  this  undergoes  a  slight 
decomposition  on  standing,  so  that  ultimately  a  solution  is  obtained 
containing  about  98  per  cent  of  the  compound  and  2  per  cent  of 
water,  and  this  is  what  is  called  pure  nitric  acid. 

Properties.  Nitric  acid  is  a  colorless  liquid.  It  has  a  density  of  1.51, 
and  boils  at  86°,  with  partial  decomposition. 

An  aqueous  solution  containing  68  per  cent  of  the  acid  has  a  con- 
stant boiling  point  (p.  127)  and  distills  with  unchanged  concentration. 
This  solution  has  a  density  of  1.4  and  constitutes  the  concentrated 
acid  of  commerce. 


1 


COMPOUNDS  OF  NITROGEN  179 

* 

Chemical  conduct  of  nitric  acid.  Some  of  the  most  important  reactions 
of  nitric  acid  are  the  following : 

1.  Acid  properties.  As  the  name  indicates,  it  is  an  acid  and  has  all 
the  properties  characteristic  of  that  class  of  substances.    When  dis- 
solved in  water  it  forms  the  ions  H+  and  NO3^    The  solution  changes 
blue  litmus  red  and  neutralizes  bases  forming  salts.    It  is  one  of  the 
strongest  of  acids  (p.  155). 

2.  Unstable  character.  The  unstable   character   of  nitric    acid  has 
already  been  described  (p.  178). 

3.  Oxidizing  action.  Since  nitric  acid  contains  a  large  percentage  of 
oxygen  and  readily  decomposes  with  evolution  of  oxygen,  it  serves  as 
a  strong  oxidizing  agent.   Under  ordinary  circumstances,  in  the  pres- 
ence of  a  substance  readily  oxidized,  the  acid  decomposes  according 
to  the  following  equation : 

2  HNO3  =  H2O  +  2  NO  4-  3  O 

In  such  cases  oxygen  is  not  evolved,  but  enters  into  combination 
with  the  oxidizable  substance  present.  In  this  way  carbon,  when 
heated  with  nitric  acid,  is  oxidized  to  carbon  dioxide.: 

C  +  2  O  =  C02 

4.  Action  upon  organic  compounds.  Nitric  acid  readily  reacts  with 
many  organic  substances,  forming  compounds  of  great  importance. 
Thus,  with  ordinary  glycerin  it  forms  the  compound  'known  as  nitro- 
glycerin,  which  is  the  explosive   constituent  of  dynamite ;  likewise, 
with  cellulose,  the  principal  constituent  of  wood-fiber,  it  forms  nitro- 
celluloses,  which  are  used  in  making  smokeless  gunpowder.    When 
nitric  acid  acts  upon  protein  matter,  a  yellow  compound  known  as 
xanthoprotein  is  formed ;  hence  nitric  acid  in  contact  with  the  skin 
produces  a  yellow  stain. 

5.  Action  upon  metals.  All  of  the  metals,  with  the  exception  of  gold, 
platinum,  and  a  few  of  the  rare  metals,  are  acted  upon  more  or  less 
readily  by  nitric  acid.    In  discussing  the  action  of  nitric  acid  upon 
these  metals  it  is  convenient  to  divide  them  into  two  classes : 

(a)  Metals  having  a  higher  electrode  potential  than  hydrogen.  It  will 
be  recalled  that  any  of  the  metals  occurring  above  hydrogen  in  the 
electromotive  series  of  the  metals  (p.  158)  will  in  general  liberate 
hydrogen  from  dilute  acids.  At  first  thought  it  might  be  expected 
that  nitric  acid  would  act  upon  these  metals  in  a  similar  way.  When 


180  GENERAL  CHEMISTRY 

one  reflects,  however,  that  nitric  acid  is  a  strong  oxidizing  agent,  while 
hydrogen  has  strong  reducing  properties,  it  seems  reasonable  to  sup- 
pose that  nitric  acid  would  be  reduced  by  the  hydrogen,  yielding  re- 
duction products.  Experiments  show  that  this  is  what  actually  happens. 
The  particular  reduction  products  formed  in  any  case  depend  upon 
the  metal,  the  concentration  of  the  acid,  and  the  conditions  under 
which  the  reaction  takes  place.  The  following  compounds  represent 
successive  steps  in  the  reduction  of  nitric  acid: 

HNO3  —  >-  NO2  —  *-  HNO2  —  H  NO 


It  is  possible,  by  selecting  different  metals  and  by  modifying  the 
conditions  of  the  reaction,  to  obtain  any  of  these  products.  Under  or- 
dinary conditions,  however,  either  nitric  oxide  Q^O),  nitrogen  dioxide 
(NO2),  or  a  mixture  of  the  two  is  generally  formed.  The  course  of  the 
reaction  may  be  sjiown  by  the  study  of  a  typical  example,  such  as  the 
action  of  nitric  acid  of  medium  concentration  upofc  zinc.  The  first  step 
in  the  reaction  consists  in  the  formation  of  zinc  nitrate  and  hydrogen  : 

Zn  +  2  HNO3  =  Zn(NO3)2  +  2  H  (1) 

The  hydrogen  is  not  evolved  as  such,  since  it  at  once  reacts  with  the 

3  H  +  HN03  =  2  H20  +  NO  (2) 

The  products  of  the  reaction  between  zinc  and  the  acid,  therefore,  are 
zinc  nitrate,  water,  and  nitric  oxide. 

It  is  often  convenient  to  express  the  reaction  in  a  single  equation.  This  is 
readily  done  by  combining  equations  (1)  and  (2)  as  given  above.  Before  the 
equations  are  combined,  however,  they  must  be  modified  so  as  to  express  the 
fact  that  all  the  hydrogen  represented  as  formed  according  to  equation  (1) 
reacts  with  the  nitric  acid  according  to  equation  (2).  This  may  be  done  by 
multiplying  the  first  equation  by  3  and  the  second  equation  by  2.  The  two 
equations  will  then  be  as  follows  : 

3  Zn  +  6  HNO3  =  3  Zn(NO3)2  +  6  H 
6  H  +  2  HNO3  =  4  H2O  +  2  NO 

By  canceling  the  common  factor  6  H,  which  represents  the  hydrogen  formed  in 
the  one  reaction  and  consumed  in  the  other,  and  then  combining  the  equations, 
the  following  is  obtained  : 

3  Zn  +  8  HNO3  =  3  Zn(NO3)2  +  4  H2O  +  2  NO 

This  complete  equation  has  the  advantage  of  making  it  possible  to  calculate  very 
easily  the  proportions  in  which  the  various  substances  enter  into  the  reaction  or 
are  formed  in  it.  It  is  unsatisfactory  in  that  it  does  not  give  full  information 


COMPOUNDS  OF  NITROGEN  181 

about  the  way  in  which  the  reaction  takes  place.  For  example,  it  does  not  sug- 
gest that  hydrogen  is  at  first  formed,  and  subsequently  transformed  into  water. 
It  is  always  much  more  important  to  remember  the  steps  in  a  chemical  reaction 
than  to  remember  the  equation  expressing  the  complete  action,  for  if  these 
steps  are  understood,  the  complete  equation  is  easily  obtained  in  the  manner 
just  described. 

(b)  Metals  having  a  lower  electrode  potential  than  hydrogen.  Those 
metals  occurring  below  hydrogen  in  the  electromotive  series,  when 
acted  upon  by  nitric  acid,  are  first  oxidized  to  the  corresponding 
oxides.  The  acid  is  thereby  decomposed  into  a  compound  having  a 
lower  percentage  of  oxygen,  such  as  nitrous  or  hyponitrous  acid,  or 
it  may  form  water  and  one  or  more  of  the  oxides  of  nitrogen.  The 
oxide  of  the  metal,  with  one  or  two  exceptions  (see  antimony,  p.  371), 
dissolves  in  the  acid,  forming  the  corresponding  nitrate  and  water. 

The  course  of  the  reaction  may  best  be  shown  by  a  study  of  some  typical 
examples.  When  moderately  dilute  nitric  acid  (density  1.2)  acts  upon  copper,  the 
reaction  may  be  expressed  by  the  following  equations : 

2  HNO3  =  H2O  +  2  NO  +  3  O 
3  O  +  3  Cu  =  3  CuO 
3  CuO  4-  6  HNO3  =  3  Cu(NO3)2  +  3  H2O 

By  canceling  the  factors  3  O  and  3  CuO,  representing  substances  .formed  in 
one  reaction  and  used  up  in  another,  and  combining  these  three  equations,  the 
following  equation  is  obtained  : 

3  Cu  +  8  HNO3  -  3  Cu(NO3)2  +  2  NO  +  4  H2O 

If  concentrated  acid  is  used  in  place  of  dilute,  nitrogen  dioxide  is  liberated,  as 
shown  in  the  following  equations  : 

2  HXO3  =  H2O  +  2  NO2  4-  O 
Cu  +  O  =  CuO 
CuO  +  2  HNO3  =  Cu(NO3)2  +  H2O 

Combining  these  into  a  single  equation,  we  obtain  the  following : 
Cu  +  4  HNO3  =  Cu(NO3)2  +  2  H2O  +  2  NO2 

Structural  formula  of  hydrogen  nitrate.  Since  hydrogen  nitrate  is 
largely  ionized  in  aqueous  solution,  it  is  evident  that  in  dealing  with 
such  solutions  we  have  to  do  not  so  much  with  the  nitrate  itself  as 
with  the  ions  which  it  forms,  namely,  H+  and  NO3~.  The  reactions  of 
the  pure  compound,  on  the  other  hand,  indicate  that  the  atoms  present 
in  the  molecule  are  combined  as  represented  in  the  following  structural 

formula :  ,  n 

H-0-N«°- 


182  GENERAL  CHEMISTRY 

Salts  of  nitric  acid ;  the  nitrates.  The  salts  of  nitric  acid  are  called 
nitrates.  They  can  be  obtained  by  the  general  method  used  for  pre- 
paring salts,  such  as  the  action  of  nitric  acid  upon  the  metals  them- 
selves, or  upon  their  oxides  or  hydroxides.  Some  of  these,  especially 
sodium  nitrate  and  potassium  nitrate,  are  found  in  nature.  The 
nitrates  of  most  of  the  metals  are  white  solids.  The  nitrate  of  copper 
is  blue,  that  of  nickel  is  green,  while  cobalt  nitrate  is  cherry-red.  All 
of  the  normal  nitrates  are  soluble  in  water  and  form  the  ions  M+  and 
NOg~,  in  which  M  represents  the  metal.  Thus,  potassium,  calcium, 
and  iron  nitrates  ionize  as  represented  by  the  following  equations : 

KNO3  +=t  K+,  NO8~ 

Ca(N03)2:«=>:Ca++,  NO,-,  NO8~ 
Fe(N03)3^=±Fe+++,  NO8%  NO,-,  NO3~ 

When  heated,  the  nitrates  undergo  decomposition.  As  a  rule,  the 
metal  is  left  in  the  form  of  the  oxide,  while  oxygen  and  oxides  of 
nitrogen  are  evolved.  Thus,  when  copper  nitrate  is  heated,  the  blue 
color  of  the  nitrate  gradually  gives  way  to  the  black  color  of  copper 
oxide,  while  the  evolution  of  nitrogen  dioxide  is  indicated  by  the 
reddish  color  of  the  evolved  gas: 

2  Cu(N03)2  =  2  CuO  +  4  NO2  +  O2 

In  the  case  of  a  few  of  the  nitrates,  however,  oxygen  alone  is  evolved. 
Thus,  sodium  nitrate,  when  heated,  forms  sodium  nitrite  (NaNO2) 
and  oxygen:  2  NaNO8  =  2  NaNO2  +  O2 

The  nitrates  have  various  uses.  Large  quantities  of  sodium  nitrate 
are  exported  from  Chile  to  different  countries,  where  it  is  used  in  the 
preparation  of  nitric  acid  and  in  the  manufacture  of  fertilizers.  In 
European  countries  calcium  nitrate  (Ca(NO3)2),  prepared  by  neutral- 
izing the  dilute  solution  of  nitric  acid  obtained  by  the  Berkeland  and 
Eyde  process  (p.  177)  with  lime  (CaO),  is  coming  to  be  used  as  a 
fertilizer,  under  the  name  air  saltpeter. 

Nascent  state.  We  have  seen  that  when  nitric  acid  acts  upon  metals 
having  a  higher  electrode  potential  than  hydrogen,  the  hydrogen  is 
not  evolved  as  such,  but  reacts  with  the  nitric  acid  present,  forming 
various  reduction  products.  If  hydrogen  is  generated  in  a  separate 
flask,  however,  and  the  pure  gas  is  conducted  into  the  acid  at  ordinary 
temperatures,  no  such  reduction  takes  place.  Evidently  this  differ- 
ence in  the  action  of  the  hydrogen  is  connected  with  the  fact  that  in 


COMPOUNDS  OF  NITROGEN  183 

the  one  case  the  hydrogen  is  generated  in  contact  with  the  nitric  acid, 
while  in  the  other  case  it  is  not.  This  is  but  one  example  of  many 
reactions  which  seem  to  indicate  that  the  activity  of  an  element  is 
greatest  just  at  the  instant  of  its  liberation  from  its  compounds. 
Elements  in  this  condition  are  said  to  be  in  the  nascent  state,  the 
word  nascent  being  derived  from  a  Latin  word  meaning  "  to  be  born." 
This  greater  activity  is  usually  explained  upon  the  assumption  that 
an  element  at  the  instant  of  liberation  from  its  compound  is  in  the 
form  of  atoms,  and  is  therefore  more  reactive  than  after  the  atoms 
have  combined  to  form  molecules  (see  character  of  compounds  of 
nitrogen,  p.  167).  The  increased  activity  has  also  been  ascribed  to 
various  other  causes,  such  as  to  the  energy  liberated  when  an  element 
is  set  free  from  compounds,  as  well  as  to  the  catalytic  action  of  the 
substances  in  contact  with  the  element  when  liberated. 

Nitrous  acid  (HN02).  While  sodium  nitrate  may  be  decomposed  by 
heat  into  sodium  nitrite  and  oxygen,  a  high  temperature  is  required 
to  effect  the  change.  The  formation  of  the  nitrite  takes  place  much 
more  readily  if  the  nitrate  is  mixed  with  a  mild  reducing  agent,  such 
as  lead,  before  it  is  heated : 

NaNO3  +  Pb  =  PbO  +  NaNO2 

On  treating  the  resulting  mass  with  water,  the  nitrite  dissolves  and 
may  be  filtered  from  the  insoluble  lead  oxide.  When  the  resulting 
solution  is  evaporated,  the  nitrite  is  obtained  in  the  form  of  a  white 
solid.  This  compound  is  the  sodium  salt  of  nitrous  acid.  When  this 
is  treated  with  sulfuric  acid,  therefore,  we  should  expect  to  have 
nitrous  acid  liberated: 

NaNO2  +  H2SO4  =  NaHSO4  +  HNO2 

When  the  reaction  is  carried  out,  however,  we  obtain  not  nitrous  acid, 
but  a  mixture  of  nitric  oxide  and  nitrogen  dioxide.  It  is  probable 
that  nitrous  acid  is  at  first  formed,  but,  being  unstable,  it  decomposes 
at  once  into  water  and  nitrogen  trioxide  (N2O3),  the  latter  compound 
in  turn  yielding  nitric  oxide  and  nitrogen  dioxide.  These  reactions 
are  expressed  in  the  following  equations : 


N2O8  =  NO  +  NO2 

It  is  evident,  therefore,  that  nitrous  acid  is  very  unstable.    When  a 
mixture  of  nitric  oxide  and  nitrogen  dioxide  is  passed  into  cold  water, 

/U>— J^    ° 


184  GENERAL   CHEMISTRY 

the  resulting  liquid  contains  a  slight  amount  of  nitrous  acid,  but  this 
soon  decomposes,  forming  nitric  acid,  as  represented  in  the  following 
equation  :  g  HN(  =  HNQ  +  2  NO  +  H 


Inasmuch  as  nitrous  acid  combines  with  oxygen,  forming  nitric  acid, 
it  acts  as  a  reducing  agent  in  the  presence  of  substances  that  readily 
yield  oxygen.  On  the  other  hand,  in  the  presence  of  a  substance  hav- 
ing a  strong  affinity  for  oxygen,  nitrous  acid  acts  as  an  oxidizing  agent, 
decomposing  as  represented  in  the  following  equation  : 

2  HNO2  -  H2O  +  2  NO  +  O 

Salts  of  nitrous  acid  ;  the  nitrites.  While  nitrous  acid  itself  is  very 
unstable,  its  salts,  the  nitrites,  are  comparatively  stable.  Like  the 
nitrates,  most  of  these  salts  are  solids  and  soluble  in  water.  The 
nitrites  of  potassium  and  of  sodium  are  common  reagents  in  the 
chemical  laboratory. 

Hyponitrous  acid.  Just  as  nitric  acid,  under  the  influence  of  reducing  agents, 
gives  up  oxygen,  forming  nitrous  acid,  so  the  latter  compound,  when  similarly 
treated,  loses  oxygen,  forming  a  compound  known  as  hyponitrous  acid.  This 
compound  contains  hydrogen,  nitrogen,  and  oxygen  in  the  proportion  represented 
by  the  formula  HXO.  The  determination  of  the  molecular  weight  of  the  com- 
pound, however,  shows  that  it  is  just  double  that  represented  by  the  formula 
HNO.  Since  the  formula  is  always  intended  to  represent  the  composition  of  the 
molecule,  we  must  assign  to  hyponitrous  acid  the  formula  H2N2O2.  This  acid, 
while  very  unstable,  has  nevertheless  been  obtained  in  the  form  of  white  crys- 
tals. These  dissolve  in  water,  forming  a  solution  of  the  acid,  which,  however, 
soon  decomposes  into  water  and  nitrous  oxide': 

H2N2O2  =  H2O  +  N2O 

The  salts  of  the  acid  are  known  as  hyponitrites.  A  number  of  these  have  been 
prepared.  The  silver  salt  is  a  yellow  solid  rather  easily  obtained,  since  it  is  only 
slightly  soluble  in  water. 

Hydroxylamine  (NH2OH).  This  compound  may  be  regarded  as  ammonia  in  which 
an  atom  of  hydrogen  has  been  replaced  by  the  radical  OH.  The  name  hydroxyl- 
amine  indicates  the  presence  of  the  hydroxyl  group,  as  well  as  of  the  amino 
group  (NH2).  Hydroxylamine  was  first  obtained  in  dilute  solution  in  1865, 
but  the  compound  was  not  isolated  in  a  pure  state  until  1891,  when  de  Bruyn 
obtained  it  in  the  form  of  white  needles  melting  at  33°.  It  is  formed  by  the 
action  of  hydrogen  on  nitric  acid  : 

HN03  +  3  H2  =  NH2OH  +  2  H2O 

Hydroxylamine  resembles  ammonia  in  being  soluble  in  water,  combining  with  it 
to  form  the  base  (NH3OH)  OH.  This  base,  however,  is  much  less  ionized  and 
therefore  weaker  than  the  corresponding  ammonium  compound.  By  the  action 


COMPOUNDS  OF  NITROGEN  185 

of  strong  acids  the  corresponding  salts  may  be  obtained.  Thus,  with  hydrochloric 
acid  there  is  formed  the  chloride  (NH3OH)C1.  These  salts  are  much  more  stable 
than  the  free  body  and  are  therefore  more  largely  used.  It  is  a  strong  reducing 
agent,  its  reducing  properties  being  much  more  marked  than  those  of  ammonia. 
Its  chief  importance  lies  in  its  reactions  with  certain  types  of  carbon  compounds. 


COMPOUNDS  OF  NITROGEN  AND  OXYGEN 
The  names  and  formulas  of  the  oxides  of  nitrogen  are  as  follows  : 
Nitrous  oxide  (hyponitrous  anhydride)  (N2O):  a  colorless  gas. 
Nitric  oxide  (NO)  :  a  colorless  gas. 

Nitrogen  dioxide  (NO2)  :  a  gas,  deep  reddish-brown  in  color. 
Nitrogen  trioxide  (nitrous  anhydride)  (^Og):   exists  only  at  low 
temperatures,  both  in  liquid  and  solid  form. 

Nitrogen  tetroxide  (N2O4)  :   a  low-boiling,  nearly  colorless  liquid. 
Nitrogen  pentoxide  (N2O5)  :  a  white,  crystalline  solid. 
Nitrogen  hexoxide  (N2O6)  :  an  unstable  greenish  solid. 

Nitrous  oxide  (N20).  Nitrous  oxide  was  first  prepared  by  Priestley 
in  1772.  Davy  determined  its  composition  in  1800  and  was  the  first 
to  point  out  the  property  which  the  gas  possesses  of  rendering  one 
temporarily  unconscious  when  it  is  inhaled. 

Preparation.  Nitrous  oxide  can  be  prepared  by  the  action  of  suitable 
reducing  agents  upon  nitric  acid,  as  well  as  upon  nitric  oxide.  Priestley 
first  obtained  it  from  this  latter  compound  by  the  reducing  action  of  iron  * 


The  most  convenient  method  for  its  preparation  consists  in  heating 
ammonium  nitrate.  Just  as  ammonium  nitrite,  when  heated,  yields 
water  and  nitrogen  (p.  106),  so  ammonium  nitrate  decomposes  in  a 
similar  way,  forming  water  and  nitrous  oxide.  The  similarity  between 
the  two  reactions  is  shown  in  the  following  equations  : 


The  decomposition  is  effected  by  heating  the  ammonium  nitrate  in  a 
flask  such  as  that  used  in  the  preparation  of  oxygen  (Fig.  6,  p.  19). 
The  nitrous  oxide  is  evolved  and  may  be  collected  over  water  — 
preferably  warm,  since  the  gas  is  noticeably  soluble  in  cold  water. 

Properties.  Nitrous  oxide  is  a  colorless  gas  and  has  a  faint  odor. 
Its  solution  has  a  slightly  sweetish  taste.    It  is  1.53  times  as  heavy  as 


186  GENERAL  CHEMISTRY 

air,  1  1.  of  it  weighing  1.9777  g.  Its  critical  temperature  is  38.8°  and 
its  critical  pressure  7.5  atmospheres.  Liquid  nitrous  oxide  is  colorless 
and  boils  at  -  89.4°. 

Chemically,  nitrous  oxide  is  characterized  by  the  ease  with  which 
it  decomposes  into  free  nitrogen  and  oxygen.  It  is  therefore  a  good 
oxidizing  agent.  Such  substances  as  carbon,  iron,  and  phosphorus, 
when  ignited  and  introduced  into  jars  of  the  gas,  burn  with  brilliancy, 
forming  oxides  and  nitrogen.  When  inhaled,  it  produces  a  kind  of 
hysteria  (hence  the  name  laughing  gas)  and  even  unconsciousness  and 
insensibility  to  pain.  It  has  long  been  used  as  an  anesthetic  for  minor 
surgical  operations,  such  as  those  of  dentistry. 

Nitric  oxide  (NO).  This  gas  was  discovered  by  Van  Helmont  and 
was  used  by  Priestley  in  the  analysis  of  air  (p.  120). 

Preparation.  Nitric  oxide  is  most  readily  prepared  by  the  action  of 
nitric  acid  (density,  1.2)  upon  certain  metals,  such  as  copper  (p.  181). 

The  metal  is  placed  in  a  flask  A  (Fig.  75)  and  the  acid  slowly  added  through 
the  funnel  tube  B.  The  gas  escapes  through  C  and  is  collected  over  water.  The 

gas  at  first  evolved    combines  with 

B  the  oxygen  of  the  air  contained  in 

the  flask  to  form  the  reddish-brown 
nitrogen  dioxide,  but  this  is  absorbed 
as  it  bubbles  through  the  water. 


Properties.  Nitric  oxide  is  a 
colorless  gas  slightly  heavier 
than  air.  It  is  very  much  more 
difficult  to  liquefy  than  nitrous 
oxide.  Its  critical  temperature  is 
FlG  75  —  93.5°  and  its  critical  pressure 

71.2   atmospheres.     The    liquid 

boils  at  —  153°.  Nitric  oxide  is  a  much  more  stable  compound  than 
nitrous  oxide ;  nevertheless,  it  can  be  decomposed  into  its  elements 
without  difficulty.  If  a  bit  of  phosphorus  is  barely  ignited  and  at 
once  introduced  into  a  jar  of  the  gas,  the  flame  is  extinguished.  On 
the  other  hand,  if  the  phosphorus  is  first  heated  until  vigorous  com- 
bustion ensues,  and  is  then  introduced  into  the  gas,  the  combustion 
continues  with  great  brilliancy. 

One  of  the  most  characteristic  properties  of  nitric  oxide  is  its  con- 
duct toward  oxygen.  When  brought  in  contact  with  oxygen  or  air,  it 
forms  nitrogen  dioxide  (NO2). 


COMPOUNDS  OF  NITROGEN 


187 


Nitrogen  dioxide  (N02)  and  nitrogen  tetroxide  (NgOJ.  Since  these  two 
compounds  are  very  closely  related,  it  is  convenient  to  discuss  them 
together.  We  have  seen  that  nitrogen  dioxide  is  formed  by  heating 
copper  nitrate  (p.  182),  as  well  as  by  the  union  of  nitric  oxide  and 
oxygen.  Either  of  these  reactions  serves  as  a  convenient  method  for 
its  preparation.  When  the  reddish-brown  gas  so  obtained  is  cooled,  the 
color  gradually  fades  and  a  faintly  yellow  liquid  forms,  which  boils 
at  about  26°.  On  further  cooling,  the  liquid  condenses  to  an  almost 
colorless  solid  melting  at  —  10°.  This  solid  is  pure  nitrogen  tetroxide 
(N2O4).  As  the  temperature  rises,  the  reverse  change  takes  place,  the 
tetroxide  gradually  changing  back  into  the  deeply  colored  gaseous 
dioxide.  At  150°  the  change  into  the  dioxide  is  complete.  Between 
these  extremes  of  temperature  the  gas  consists  of  a  mixture  of  the 
two  compounds  in  equilibrium  with  each  other : 

N2O4^=±2NO2 

The  mixture  is  generally  referred  to,  however,  as  nitrogen  dioxide  or 
nitrogen  peroxide.    At  still  higher  temperatures  the  dioxide  is  broken 
up  into  nitric  oxide  and  oxygen.  From  what  has  been  said  it  is  evident 
that  whenever  either  the  dioxide  or  the 
tetroxide  is  generated  at  ordinary  temper- 
atures,  a   portion   of  the  one   compound 
immediately  changes   into    the    other,  so 
that  we   are  really  dealing  with  a  mix- 
ture of  the  two. 

Compounds  related  to  each  other  as 
nitrogen  dioxide  is  related  to  nitrogen 
tetroxide  are  termed  polymeric  compounds, 
or  simply  polymers.  Such  compounds  have 

the  same  percentage  composition  but  dif-     

ferent  molecular  weights. 


FIG.  76 


The  formation  of  nitrogen  dioxide  from  nitric  oxide  and  oxygen,  together 
with  the  reaction  between  the  dioxide  and  water,  may  be  shown  as  follows :  A 
tube  is  filled  with  water  and  inverted  in  a  vessel  of  water,  as  shown  in  Fig.  76. 
The  tube  is  then  nearly  filled  with  nitric  oxide,  after  which  oxygen  is  admitted, 
a  few  bubbles  at  a  time.  As  each  bubble  enters  the  tube,  the  deep  reddish-brown 
dioxide  forms.  After  a  few  moments  the  color  fades,  owing  to  the  fact  that  the 
dioxide  reacts  with  the  water,  which  at  the  same  time  rises  in  the  tube  to  take 
the  place  of  the  gas  so  removed.  It  will  be  recalled  (p.  120)  that  by  a  similar 
method  Priestley  originally  determined  what  he  called  the  "goodness  of  air." 


188  GENERAL  CHEMISTRY 

The  reaction  which  takes  place  between  the  dioxide  and  water  de- 
pends upon  the  temperature.  In  cold  water  both  nitric  and  nitrous 
acids  form:  2  NO,  +  H,O  =  HNO,  +  HNO, 

m     '.'.''  m  23 

At  a  higher  temperature  the  nitrous  acid  decomposes  as  .fast  as  formed, 
and  the  reaction  may  be  represented  as  follows  : 

3  NO2  +  H2O  =  2  HNO3  +  NO 

Nitrogen  trioxide  (N203).  When  a  mixture  of  equal  volumes  of  nitric 
oxide  and  nitrogen  dioxide  is  cooled',  there  is  obtained  at  first  a  blue 
liquid,  which,  at  a  still  lower  temperature,  condenses  to  a  solid.  The 
liquid,  as  well  as  the  solid,  is  generally  regarded  as  nitrogen  trioxide  : 


Nitrogen  pentoxide  (N205).  This  oxide  is  formed  when  nitric  acid  is 
treated  with  a  strong  dehydrating  agent,  such  as  phosphorus  pentoxide  : 

2HN03  =  H20+N205 

Nitrogfin  pentoxide  is  a  white  crystalline  solid  melting  at  29.5°.    It 
readily  decomposes  into  nitrogen  dioxide  and  oxygen  : 


It  combines  with  water,  forming  nitric  acid  : 


Anhydrides.  From  the  statements  made  it  will  be  seen  that  each  of 
the  oxygen  acids  of  nitrogen  decomposes  into  water  and  an  oxide  of 
nitrogen,  as  represented  in  the  following  equations  : 


H2N202  = 


In  the  case  of  hyponitrous  acid  and  nitrous  acid  this  decomposition 
takes  place  spontaneously,  while  with  nitric  acid  a  dehydrating  agent 
is  necessary  to  bring  about  the  reaction.  Nearly  all  acids  containing 
oxygen  undergo  a  similar  reaction,  and  the  oxides  resulting  are  termed 
anhydrides.  Thus,  nitrogen  pentoxide  is  the  anhydride  of  nitric  acid 
and  is  often  called  nitric  anhydride.  Similarly,  N2O  and  N2O3  are 
the  anhydrides  of  hyponitrous  acid  and  nitrous  acid  respectively.  The 
decomposition  of  an  acid  into  water  and  the  corresponding  anhydride 


COMPOUNDS  OF  NITROGEN 


189 


Dead  Matter 
D- 


is,  as  a  rule,  a  reversible  reaction.  Thus,  the  anhydrides  of  nitrous 
acid  and  nitric  acid  unite  with  water  to  form  the  corresponding  acids. 
The  decomposition  of  nitrogenous  organic  matter.  The  protein  matter 
occurring  in  plants  and  animals  may  be  decomposed  in  a  number  of 
ways,  the  most  important  of  which  are  the  following: 

1.  By  the  action  of  microorganisms.    Experiments  have  shown  that 
organic   matter  will   decay   only   in   the  presence   of  certain  micro- 
organisms, which  in  some  way  assist  in  its  decomposition.    In  this 
process   a  part  of  the  nitrogen  is  evolved  as  free  nitrogen,  while 
other  portions  are  converted  into  ammonia,  nitrites,  or  nitrates. 

2.  By  the  action  of  heat.    When  heated  in  the  absence  of  air,  protein 
matter   undergoes   complicated   changes,  in  which  a  portion   of  the 
nitrogen  is  evolved  in  the  form  of  ammonia. 

3.  By  the  action  of  acids  or  bases.    When  heated  to  a  high  tempera- 
ture with  concentrated  sulfuric  acid,  the  nitrogen  present  in  organic 
matter  is  converted  into  ammonia.    Since  the  ammonia  formed  in  this 
process  can  readily  be  estimated,  the  reaction  is  used  to  ascertain  the 
amount  of  nitrogen  in  organic 

matter,  such  as  foods  and  fer- 
tilizers. This  way  of  deter- 
mining nitrogen  is  known  as 
the  Kjeldahl  method.  Strong 
alkalies  have  a  similar  action 
on  nitrogen  products,  the  ni- 
trogen being  evolved  in  the 
form  of  ammonia. 

The  cycle  of  nitrogen.  The 
nitrogen  present  in  the  soil  in 
the  form  of  nitrates  is  taken 
up  by  plants  and  built  into 
protein  matter.  This  matter 
undergoes  a  series  of  changes 

which  finally  result  in  restoring  to  the  soil  the  nitrogenous  matter  orig- 
inally extracted  from  it.  The  cycle  through  which  the  nitrogen  passes 
may  be  illustrated  in  a  general  way  by  the  above  diagram  (Fig.  77). 

The  nitrogen  in  the  nitrates  (A)  of  the  soil  is  built  into  plant  protein 
(J5),  and  this  in  turn  into  annual  protein  (C).  During  life  the  nitro- 
gen assimilated  by  animals  is  largely  eliminated  in  the  form  of  urea 
(p.  289),  which  decomposes  into  ammonia,  as  indicated  in  the  figure. 


Animal, 
Protein. 
C- 


-Urea 


NUrogen 
E 


Nitrates 
FIG.  77 


190  GENERAL  CHEMISTRY  . 

After  death  the  tissue  (D)  undergoes  decay,  largely  changing  into 
ammonia  and  free  nitrogen  (J£).  The  latter  escapes  into  the  air,  while 
the  ammonia  is  absorbed  by  water  and  ultimately,  through  the  action 
of  certain  microorganisms,  changes  into  nitrites  (^),  and  finally  into 
nitrates  (A).  The  free  nitrogen  evolved  in  the  decay  of  the  dead  tissue 
may  be  built  up  again  into  compounds  through  the  action  of  the 
microorganisms  present  on  the  roots  of  certain  plants. 

It  will  be  recalled  that  the  presence  of  organic  matter  in  water 
indicates  that  the  water  has  been  polluted  and  is  unfit  for  sanitary 
purposes.  This  organic  matter  is  extremely  complex  in  character  and 
there  is  no  simple  way  of  determining  its  quantity  and  nature,  espe- 
cially when  it  is  present  in  small  amounts.  On  the  other  hand,  very 
delicate  tests  are  known  for  its  decomposition  products,  —  namely,  am- 
monia, nitrites,  and  nitrates,  —  and  these  compounds  are  not  likely  to  be 
present  unless  derived  from  organic  matter.  In  making  an  analysis  of 
water,  therefore,  the  percentages  of  these  products  are  determined,  and 
from  the  results  conclusions  are  drawn  as  to  the  purity  of  the  water. 


CHAPTER  XVI 

EQUILIBRIUM 

Introduction.  The  general  idea  of  equilibrium  as  a  balance  between 
two  transformations  opposing  each  other  has  been  developed  in  a 
number  of  instances  in  the  foregoing  pages.  It  is  desirable  to  bring 
these  cases  together  and  formulate  the  principles  which  have  been 
discovered  in  connection  with  equilibrium. 

Physical  equilibrium.  Many  examples  of  purely  physical  equilibrium 
have  already  been  described.  Thus,  the  freezing  point  of  a  liquid  has 
been  denned  as  the  temperature  at  which  the  liquid  and  solid  remain 
unchanged  in  contact  with  each  other,  the  rate  at  which  the  solid 
melts  being  just  balanced  by  the  rate  at  which  the  liquid  freezes. 
Vapor  pressure  at  a  given  temperature  is  the  equilibrium  value  of  the 
rate  at  which  molecules  leave  the  liquid,  as  compared  with  the  rate  of 
their  return.  Saturation  is  reached  when  the  rate  at  which  a  solid  or 
a  gas  dissolves  is  just  compensated  by  the  rate  of  its  deposit  or  escape 
from  solution. 

In  all  these  cases,  when  equilibrium  is  reached,  the  condition  is  not 
one  of  rest  but  of  motion ;  the  number  of  molecules  moving  in  one 
direction  is  just  equal  to  the  number  moving  in  the  other. 

Chemical  equilibrium.  When  chemical  action  takes  place  between 
two  substances,  it  often  happens  that  the  action  is  incomplete,  some  of 
both  materials  apparently  failing  to  take  part  in  the  reaction.  Experi- 
ment has  shown  that  in  these  cases  there  results  an  equilibrium  very 
similar  to  those  of  physical  equilibrium  just  enumerated.  For  an 
understanding  of  the  conditions  which  produce  such  an  equilibrium, 
it  will  be  necessary  first  to  make  a  study  of  the  factors  influencing  the 
speed  of  those  reactions  which  reach  completion. 

Factors  which  influence  the  speed  of  reactions.  By  the  speed  of  a 
reaction  is  meant  the  quantity  of  material  undergoing  transformation- 
in  a  unit  of  time.  Quite  a  number  of  factors  are  involved  in  determin- 
ing this  speed. 

1.  Affinity.  The  specific  attraction  between  the  reacting  substances, 
which  we  call  chemical  affinity,  is  of  fundamental  importance.  Some 

191 


192  GENERAL   CHEMISTRY 

substances,  such  as  fluorine  and  oxygen,  will  not  unite  at  all.  Under 
certain  conditions  others,  as  hydrogen  and  oxygen,  act  upon  each  other 
with  great  energy.  Other  things  being  equal,  the  greater  the  affinity 
between  the  substances  the  greater  the  speed  of  their  reaction.  The 
speed  may  therefore  be  employed  as  a  measure  of  the  affinity. 

2.  Temperature.    A   rise   in   temperature   always  greatly  increases 
the  speed  of  a  reaction.    As  a  rough  approximation  Ostwald  esti- 
mates that  a  rise  of  10°  about  doubles  the  speed.    If,  therefore,  a 
reaction  is  proceeding  at  a  given  speed  at  0°,  it  will  go  twice  as 
fast  at  10°  and  four  times  as  fast  at  20°.    At  100°  its  speed  will 
be  1024  times  as  great  as  at  0°.    Reactions  whose  speed  is  imper- 

.ceptible  at  ordinary  temperatures  (for  example,  the  oxidation  of  coal) 
may  therefore  become  very  rapid  at  a  temperature  within  the  reach  of 
a  Bunsen  burner. 

3.  Qatalysis.  The  change  in  the  speed  of  a  reaction  due  to  catalytic 
agents  has  been  commented  upon  in  several  instances,  as  in  the  de- 
composition of  potassium  chlorate  in  the  presence  of  manganese  diox- 
ide.   It  has  been  found  that  the  speed  of  the  great  majority  of  reactions 
can  be  modified  by  some  suitable  catalyzer,  and  that  it  may  be  either 
hastened  or  retarded  in  this  way.    The  effect  produced  by  an  accel- 
erating catalyzer  is  the  same  as  if  the  reaction  were  to  be  carried  on 
at  a  higher  temperature,  so  that  it  is  often  more  economical  to  employ 
a  suitable  catalyzer  than  to  be  at  the  expense  of  maintaining  the 
higher  temperature.    Since  the  catalyzer  merely  changes  the  speed  of 
reaction,  and  does  not  add  any  energy  to  the  reacting  bodies,  it  can- 
not bring  about  a  reaction  which  does  not  take  place  by  itself;  for 
example,  it  cannot  maintain  an  endothermic  reaction. 

4.  Concentration.    Since,  as  we  believe,  reactions  take  place  between 
individual  molecules,  and  since  these  cannot  act  upon  each  other  at 
a  distance,  it  will  be  seen  that  any  condition  which  increases  the  fre- 
quency of  their  meeting  will  promote  the  speed  of  reaction.    The  more 
molecules  there  are  in  a  given  space  the  faster  the  reaction  will  pro- 
ceed, so  that  the  molecular  concentration  will  greatly  affect  the  speed 
of  reaction.    In  cases  where  a  gas  is  one  of  the  reacting  bodies,  the 
pressure  under  which  the  gas  is  maintained  will  have  great  influence, 
not  because  of  any  physical  effect  of  pressure,  but  because  the  pres- 
sure determines  the  concentration  of  the  gas,  Boyle's  law  reminding 
us  that  when  the  pressure  is  doubled,  the  same  number  of  molecules 
is  crowded  into  half  the  space. 


EQUILIBRIUM  193 

5.  Solution.  Solution  promotes  the  speed  of  reaction  in  another  way, 
for  it  makes  possible  a  free  movement  of  the  molecules,  and  thus 
brings  them  into  frequent  contact. 

The  law  of  mass  action.  Of  the  five  factors  just  enumerated,  the 
-effect  of  concentration  upon  the  speed  is  most  easily  measured.  As 
early  as  the  beginning  of  the  nineteenth  century  Berthollet  made  im- 
portant measurements  of  this  kind,  but  it  was  not  until  1867  that  the 
relation  of  concentration  to  the  speed  of  a  reaction  was  definitely 
formulated.  In  this  year  the  Norwegian  scientists  Guldberg  and 
Waage  published  the  results  of  their  investigations  in  the  form  of  the 
law  of  mass  action,  which  states  that  the  speed  of  a  reaction  is  propor- 
tional to  the  molecular  concentrations  of  the  reacting  substances.  It 
must  be  kept  in  mind  that  by  molecular  concentration  is  meant  the 
number  of  gram-molecular  weights  per  liter. 

To  illustrate  the  meaning  of  this  law,  let  us  take  the  reaction  ex- 
pressed by  the  equation 

Let  [a]  and  [6]  be  the  molecular  concentrations  of  the  nitrogen  and 
oxygen  at  the  outset.  The  speed  will  be  proportional  to  the  number 
of  nitrogen  molecules  in  unit  volume  [a],  and  also  to  that  of  the 
oxygen  molecules  [&].  In  accordance  with  a  simple  principle  of 
algebra  it  will  therefore  be  proportional  to  the  product  [«]  x  [6].  It 
will,  of  course,  depend  upon  the  temperature,  the  presence  of  cata- 
lyzers, the  affinities  of  the  reacting  substances,  and  the  units  em- 
ployed in  the  measurement.  At  a  given  temperature  and  with  a  given 
set  of  units,  all  these  influences  are  constant  in  value,  and  we  may 
designate  their  combined  effect  by  the  constant  &,  which  is  called  the 
affinity  constant  at  that  temperature.  We  shall  then  have  the  equation 

speed  =  [a]  x  [6]  x  k 

In  the  formation  of  ammonia  from  its  elements,  where  more  than  one 
molecule  of  a  given  kind  enters  into  the  reaction,  the  full  equation  is 


Evidently  the  speed  is  not  proportional  merely  to  [a],  the  concen- 
tration of  the  hydrogen,  but  to  the  product  [a]  x  [a~\  x  [a]  x  [5], 
so  that  we  have  the  equation 

speed  =  [a]3  x  [6]  X  k 


194  GENERAL  CHEMISTRY 

Active  mass.  As  long  as  we  deal  with  reactions  in  gas  mixtures  or 
in  solution,  the  concentrations  of  the  reacting  bodies  are  easy  to  define. 
Such  reactions  are  called  homogeneous.  When  a  solid  is  acted  upon  by 
a  gas,  as  in  the  reaction  of  steam  upon  iron,  or  by  a  liquid,  as  when 
hydrochloric  acid  acts  upon  zinc,  the  reaction  is  said  to  be  heteroge- 
neous, because  the  materials  are  not  mixed  with  each  other  in  even 
concentration.  In  the  reaction  expressed  by  the  equation 

4H20  +  3Fe ^Fe3O4  +  4H2 

we  have  no  difficulty  in  defining  the  concentration  of  the  steam  and 
in  seeing  that  every  molecule  of  it  is  available  for  the  reaction.  With 
the  iron  the  case  is  quite  different.  Only  those  molecules  forming  the 
surface  of  the  solid,  or  given  off  as  vapor,  can  have  any  effect  upon 
the  velocity  of  the  reaction  at  a  given  time,  and  in  either  case  the 
effect  will  be  practically,  though  not  accurately,  constant.  That 
portion  of  a  substance  which  is  available  for  a  reaction  at  a  given 
time  is  called  its  active  mass,  and  it  is  with  the  concentration  of  the 
active  mass  that  the  law  of  mass  action  is  concerned.  The  best  we 
can  do  in  regard  to  the  active  mass  of  a  solid  is  to  consider  it  as  a 
constant  during  the  reaction : 

4H2O  +  3Fe >-Fe3O4  +  4H2,    or    [>]4  x  kl  X  k  =  speed 

[a]  kl 

In  this  equation  k1  is  the  constant  effect  of  the  solid  iron  and  k  the 
constant  effect  of  chemical  affinity  and  temperature.  From  this  it 

follows  that  ,       _  _ 

speed  =  [a\  x  K 

in  which  K  is  the  fourth  root  of  the  product  of  the  one  constant 
by  the  other,  and  is  therefore  a  constant  also.  This  means  that 
the  only  effect  produced  by  varying  the  quantity  of  materials  in 
this  reaction  is  due  to  the  steam,  all  other  forces  being  constant. 
The  concentration  of  the  steam  therefore  determines  the  speed  of 
the  reaction. 

Irreversible  reactions.  Many  reactions  are  known  which  go  steadily 
on  to  completion.  Thus,  potassium  chlorate  decomposes  into  potas- 
sium chloride  and  oxygen,  and  there  is  no  evidence  that  the  products 
of  the  reaction  ever  act  directly  upon  each  other  to  form  again  the 
original  compound.  In  a  similar  way  sugar  decomposes  into  carbon, 
water,  and  other  products.  The  chemical  action  in  many  explosives 
is  also  of  this  type. 


EQUILIBRIUM 


195 


Reversible  reactions.  In  the  majority  of  cases,  however,  the  reaction 
may  be  reversed  by  properly  choosing  the  conditions.  Thus,  if  steam  is 
continuously  passed  over  heated  iron,  all  of  the  iron  is  finally  converted 
intooxide:  4  H,O  +  3  Fe -+ Fe3O4  +  4  H, 

On  the  other  hand,  by  continuously  passing  hydrogen  over  heated 
oxide  the  latter  may  be  completely  reduced  to  metallic  iron: 

4  H2  +  Fe304 +  3  Fe  +  4  H2O 

In  a  similar  way  nitrogen  and  hydrogen,  when  heated  (by  an  electric 
discharge),  combine  to  form  ammonia,  and  ammonia  may  in  turn  be 
decomposed  into  its  constituent  elements. 

Equilibrium.  In  many  cases  the  conditions  favorable  to  the  two  re- 
actions approach  very  closely  to  each  other,  so  that  the  reactions  go 
on  simultaneously  in  the  same  mixture,  the  speeds  being  not  very 
different.  In  these  cases,  if  we  follow  the  speed  of  one  of  the  reac- 
tions, it  appears  to  become  slower  and 
slower,  until  finally  the  reaction  ceases 
altogether,  stopping  before  all  of  the 
available  material  has  been  used  up. 
What  really  happens  is  that  the 
reverse  reaction,  which  increases  in 
speed  as  the  concentration  of  the 
reaction  products  increases,  finally 
equals  the  speed  of  the  direct  reac- 
tion, resulting  in  a  state  of  balance, 
or  equilibrium.  The  equations 

4H20  +  3Fe:r=>:Fe304  +  4H2 
and      N2  +  3H2^=±2NH3 

represent  this  balanced  condition  and 
are  called  equilibrium  equations. 


Demonstration  of  equilibrium.    The  fact  -pIG   -g 

that  nitrogen  and  hydrogen  reach  an  equi- 
librium with  ammonia  may  be  demonstrated  very  easily.  Approximately  3 
volumes  of  hydrogen  and  1  of  nitrogen  are  inclosed  over  mercury  in  a  eudi- 
ometer tube  A  (Fig.  78).  Electric  sparks  from  a  small  induction  coil  B  are 
passed  through  the  mixture  for  some  time.  At  first  there  is  a  slow  contraction 
in  volume.  When  this  ceases,  a  little  sulf  uric  acid  is  introduced  over  the  mercury 
through  the  stopcock  C.  The  acid  absorbs  the  ammonia  formed  in  the  reaction. 
The  volume  formed  can  thus  be  readily  determined. 


196  GENERAL  CHEMISTRY 

If  ammonia  gas  is  inclosed  in  a  second  eudiometer,  and  sparks  are  passed  for 
quite  a  time,  a  steady  expansion  in  volume  occurs,  owing  to  the  formation  of 
nitrogen  and  hydrogen  from  the  ammonia.  When  this  expansion  ceases,  the 
volume  of  the  remaining  ammonia  may  be  determined  as  before,  by  introducing 
a  little  sulf uric  acid.  When  the  experiment  is  carefully  conducted,  the  percentage 
of  the  ammoriia  by  volume  is  the  same  in  each  case,  showing  that  the  same  equilib- 
rium is  reached  whether  the  reaction  is  conducted  in  one  direction  or  the  other. 

If  sulfuric  acid  is  added  at  the  outset  of  the  first  experiment,  the  contraction 
continues  as  long  as  the  sparks  are  passed  through  the  mixture,  for  the  ammonia 
is  removed  as  fast  as  it  is  formed,  and  no  equilibrium  can  be  reached. 

Factors  affecting  the  point  of  equilibrium.  Since  the  point  reached 
in  equilibrium  is  merely  the  balance  between  the  speeds  of  two  oppos- 
ing reactions,  any  of  the  factors  which  affect  the  speed  of  either  reac- 
tion taken  separately  will  affect  the  point  of  equilibrium.  Thus,  the 
affinities  in  the  two  reactions  are  in  general  different,  tending  to  throw 
the  point  of  equilibrium  toward  one  or  the  other  extreme.  In  general, 
change  of  temperature  does  not  affect  the  velocity  of  the  two  reactions 
to  the  same  extent,  and  so  results  in  a  shift  of  the  equilibrium  point. 
As  far  as  is  known,  a  catalyzer  affects  both  equally.  It  therefore  has 
no  effect  upon  the  point  of  equilibrium,  but  merely  hastens  the  reach- 
ing of  equilibrium.  The  relative  concentvrp.tinris  of  the  rearting-kad^s 
have,  as  we  have  seen,  a  great  influence  upon  the  speed  of  a  reaction, 
and  consequently  upon  the  equilibrium  point.  This  effect  must  now 
be  considered  more  in  detail. 

Equilibrium  equations.  Let  us  suppose  that  sodium  nitrate,  sulfuric 
acid,  sodium  hydrogen  sulfate,  and  nitric  acid  are  brought  together  in 
solution,  and  that  an  equilibrium  results  as  follows : 

NaNO3  +  H2SO4  +=t  NaHSO4  +  HNO3 

[«]  [&]  [C]  [rf] 

[a],  [6],  [<?],  [cT]  being  the  respective  concentrations  at  equilibrium. 
The  speed  of  the  reaction  of  sodium  nitrate  upon  sulfuric  acid  is  pro- 
portional to  [a],  to  [5],  and  to  the  affinity  constant  of  the  reaction,  k. 

speed  =  [a]  x  [5]  x  k 

The  speed  of  the  opposing  reaction  is  expressed  in  the  equation 
speed  =  [e]  x  [d~]  x  k' 

in  which  k'  is  the  affinity  constant  for  the  action  of  nitric  acid  upon 
sodium  hydrogen-  sulfate.  At  equilibrium  these  speeds  are  equal,  and 

r  -.       r7n       ,       r  n       r  7-.       7/  \a\  x  [7>]      k' 

[«]x[J]x*=[>]x[«*]x*',    or          -        = 


EQUILIBRIUM  197 

Since  k  and  k'  are  both  constants,  their  quotient  is  also  a  constant  =  K. 
So  we  have  [a]  x  [6]  = 

M  *  [<*]  ~ 

The  constant  K,  which  is  the  ratio  of  the  two  affinity  constants,  is 
called  the  equilibrium  constant.  It  must  be  kept  in  mind  that  its  value 
is  not  at  all  dependent  on  the  concentrations  of  the  reacting  substances. 
No  matter  how  these  c^££g^00i!!ns  are  varied,  reaction  will  take  place 
in  one  direction  or  the  other  until  the  concentrations  adjust  them- 
selves to  the  value  required  by  the  equilibrium  constant. 

Types  of  equilibrium  reactions.  It  will  be  instructive  to  apply  these 
principles  to  a  few  of  the  more  common  types  of  reactions,  such  as 
will  be  met  repeatedly  in  the  following  pages. 

1.  The  reduction  of  oxides  by  hydrogen.  If  iron  oxide  and  hydrogen 
are  heated  together  in  a  sealed  tube,  an  equilibrium  results  : 

4  H2  +  Fe304^=±  3  Fe  +  4  H2O 
In  general  we  would  have  the  concentration  equation 

[H2]*  x  [Fe304] 
[H20]*  x  [Fe]3 

But  both  iron  oxide  and  iron  are  solids,  and  their  active  masses  are  con- 
stant.   Therefore 

' 


This  equation  tells  us  that  the  condition  for  equilibrium  is  that  the 
concentrations  of  the  hydrogen  and  steam  shall  be  in"  some  fixed  ratio 
to  each  other,  and  experiment  shows  that  if  the  partial  pressure  of  the 
steam  at  440°  is  10.1,  that  of  hydrogen  is  57.9,  the  ratio  being  0.174 
(=  K').  If  a  mixture  of  steam  and  hydrogen  in  this  ratio  is  passed 
over  heated  iron  or  magnetic  oxide,  or  both,  at  440°,  no  reaction  is 
produced,  and  the  relative  weights  of  the  solids  are  of  no  impor- 
tance. The  condition  for  the  complete  reduction  of  the  iron  oxide 
is  that  the  steam  shall  be  removed  as  fast  as  it  is  formed,  to  prevent 
the  reverse  reaction.  This  is  secured  when  a  current  of  hydrogen  is 
conducted  over  the  heated  oxide. 

2.  Dissociation  by  heat.  A  second  type  of  reaction  is  that  in  which 
a  substance  is  dissociated  by  heat.  A  familiar  case  is  that  of  hydrated 
salts,  a  typical  equilibrium  being  represented  in  the  equation 

CaCl2  +  2  H2O  +=±  CaCl2  •  2  H2O 


198  GENERAL  CHEMISTRY 

This  gives  the  concentration  equation 

[CaCLJ  x  [H20]2  x  k  =  [CaCl2  •  2  H2O]  x  kf 

Remembering  that  both  CaCl2  and  CaCl2  •  2  H2O  are  solids,  we  have 
[H20]2  =  JT,    and    [H2O]  -  Vjf 

At  a  given  temperature  the  equilibrium  between  the  anhydrous 
salt,  the  hydrate,  and  water  vapor  depends  entirely  upon  the  con- 
centration of  the  water  vapor,  that  is,  upon  the  vapor  pressure. 
If  the  vapor  pressure  of  the  water  in  the  atmosphere  is  greater 
than  that  required  for  equilibrium,  the  anhydrous  salt  will  keep 
absorbing  moisture,  to  reduce  it  to  that  value;  if  less,  the  hy- 
drated  salt  will  give  up  moisture,  to  bring  about  the  equilibrium. 
In  the  latter  case,  if  the  salt  is  exposed  to  the  open  air,  where 
the  moisture  is  constantly  removed  by  diffusion,  the  hydrate  will 
continuously  lose  water  and  is  said  to  be  efflorescent.  Efflorescence  is 
therefore  relative  to  the  atmospheric  humidity,  and  many  salts  would 
be  efflorescent  in  a  desert  country  which  are  not  so  in  an  ordinary 
locality. 

3.  lonization  of  an  electrolyte.  The  equations  which  express  the 
ionization  of  an  electrolyte  are  quite  similar  to  those  of  dissociation 
by  heat,  the  difference  being  that  in  the  case  of  ionization  the  equi- 
librium occurs  in  solution,  so  that  there  are  no  solids  to  be  takert 
into  account.  With  nitric  acid  the  ionization  is  expressed  by  the 

equati°n  T 

HNO.=j=tH++NO.-,    or    E 


Similarly,  for  ammonium  hydroxide  the  equilibrium  equation  is 


or 


For  strongly  ionized  substances,  such  as  strong  acids  and  bases,  as 
well  as  nearly  all  salts,  the  value  of  K  is  large,  since  even  in  nor- 
mal solution  these  are  largely  ionized  (p.  155).  For  weak  electro- 
lytes, such  as  ammonium  hydroxide,  the  value  of  K  is  a  very  small 
fractional  number. 

Effect  of  dilution  on  ionization.  -A  study  of  these  equations  will 
show  that  the  effect  of  dilution  should  be  to  increase  ionizatioi^TQ, 
make  this  clear,  let  us  suppose  that  we  start  with  a  solution  containing 


EQUILIBRIUM  199 

one  gram-molecule  of  the  electrolyte  AB  per  liter,  and  that  one  half 
of  it  is  ionized  at  this  concentration.    Then 


Now  suppose  that  a  liter  of  water  is  added.  At  first  the  concentra- 
tion of  each  ion  and  of  the  molecular  portion  of  the  electrolyte  will 
be  reduced  to  0.25  normal,  and  we  shall  have 


This  gives  a  value  for  K  much  less  than  the  original  value  (0.5),  so 
that  the  numerator  of  the  equation  must  increase  and  the  denominator 
diminish  until  the  original  value  is  restored.  This  means  that  more 
of  the  molecules  must  pass  into  the  ionic  form  until  the  quotient 
reaches  the  original  value  (0.5). 

This  effect  is  just  what  would  be  expected  from  a  kinetic  point  of 
view.  Concentration  or  dilution  will  have  little  effect  upon  the  rate 
at  which  molecules  dissociate  into  ions,  but  the  combination  of  ions 
into  molecules  will  depend  upon  how  often  they  chance  to  meet,  and 
the  more  they  are  scattered  through  a  large  volume  of  solvent  the 
less  frequently  will  this  occur. 

Strong  electrolytes  and  the  law  of  mass  action.  The  equilibrium 
constant  for  weak  electrolytes,  such  as  ammonium  hydroxide,  remains 
unchanged  over  a  wide  range  of  concentration,  the  law  of  mass  action 
applying  very  perfectly.  For  reasons  which  are  not  as  yet  entirely 
clear,  this  is  not  so  with  strongly  ionized  electrolytes,  the  value  of  the 
equilibrium  constant  changing  very  considerably  with  the  concentra- 
tion. In  all  such  cases,  as,  for  example,  with  nitric  acid,  the  equations 
just  explained  will  not  apply  with  accuracy  over  any  considerable 
range  of  concentration,  and  exact  calculations  cannot  be  based  upon 
them.  As  long  as  a  merely  qualitative  use  is  made  of  them,  however, 
they  can  be  employed  without  error,  and  it  is  in  this  way  only  that 
we  shall  have  occasion  to  make  use  of  them  in  our  reasoning. 

Conditions  for  the  completion  of  ionic  reactions.  In  the  preparation 
of  chemical  substances,  both  in  the  laboratory  and  in  the  industries, 
we  are  usually  interested  in  knowing  under  what  conditions  a  reaction 
may  be  pushed  as  far  as  possible  toward  completion.  Many  of  these 
reactions  are  carried  out  in  solution,  and  it  becomes  of  importance  to 


200  GENERAL   CHEMISTRY 

inquire  into  the  conditions  favoring  completion  of  ionic  reactions. 
There  are  three  general  conditions  which  lead  to  completed  reactions  : 
1.  The  formation  of  an  insoluble  gas.  When  we  bring  together  sodium 
nitrate  and  sulfuric  acid  in  solution,  we  have  at  the  outset  the  two 
equilibria  expressed  in  the  equations 


But  since  any  positive  ion  is  free  to  combine  with  any  negative  ion, 
the  two  ions  H+  and  NOg~  will  also  come  to  equilibrium  : 


When  the  solution  is  a  concentrated  one,  as  in  the  preparation  of 
nitric  acid,  the  equilibrium  is  reached  when  most  of  the  ions  are  com- 
bined in  the  molecular  form  HNOQ.  If  now  the  solution  is  heated 

o 

above  the  boiling  point  of  the  nitric  acid,  the  latter  distills  off  as  fast 
as  it  is  formed.  More  of  the  ions  H+  and  NO  ~  then  combine  to  form 

• 

molecular  HNOg,  in  order  to  maintain.  the  value  of  the  equilibrium 
constant  K.  This  process  continues  until  all  the  HNO3  which  can 
be  formed  has  distilled  off,  leaving  the  ions  Na+  and  HSO4~. 

The  preparation  of  acids  from  their  salts.  In  the  preparation  of  many 
acids  from  their  salts  advantage  is  taken  of  the  principle  just  de- 
scribed. Most  of  the  common  acids,  with  the  exception  of  those  of 
phosphorus,  have  boiling  points  which  are  lower  than  that  of  sulfuric 
acid  (338°).  Consequently,  when  their  salts  are  treated  with  sul- 
furic acid  and  the  resulting  liquid  warmed,  if  necessary,  the  lower- 
boiling  acid  distills  off  first.  This  method  of  preparation  is  the  one 
employed  in  all  cases  in  which  sulfuric  acid  has  no  chemical  action 
upon  the  acid  to  be  prepared. 

2.  The  formation  of  a  sparingly  soluble  solid.  When  hydrogen  chlo- 
ride (HC1)  and  silver  nitrate  (AgNO3)  are  brought  into  solution, 
the  two  equilibria  result,  as  expressed  in  the  following  equations: 


But  the  ions  Ag+  and  Cl~  will  also  form  an  equilibrium  : 


„ 


EQUILIBKIUM  201 

Now  silver  chloride  (AgCl)  is  very  difficultly  soluble  in  water,  1  1. 
dissolving  only  about  0.0016  g.,  yet  the  value  of  K  is  large,  as  is  true 
of  practically  all  salts.  The  quantity  of  silver  chloride  formed  in 
establishing  the  equilibrium  very  quickly  saturates  the  solution,  and 
the  excess  begins  to  precipitate.  This  will  continue  until  the  several 
concentrations  adjust  themselves  to  the  value  demanded  by  K.  Ob- 
viously this  will  be  when  the  concentration  of  the  ions  Ag+  and  Cl~ 
is  very  small,  practically  all  of  the  silver  having  been  precipitated 
in  the  form  of  silver  chloride.  This  leaves  the  two  ions  H+  and  NO3~ 
alone  in  solution  to  adjust  their  own  equilibrium,  so  that  the  reaction 
as  expressed  in  the  equation 

AgNO3+  HC1  =  AgCl  +  HNO3 

has  become  practically  complete  because  of  the  insolubility  of  one  of 
the  possible  equilibrium  products. 

3.  The  formation  of  very  sparingly  ionized  molecules.  When  the  base 
KOH  and  the  acid  HC1  are  brought  into  solution,  we  have  the 
equations  KOH  +=±  K+  +  OH~ 

HC1  +=±  H+  +  Cl- 

But  the  ions  H+  and  OH~  will  occasionally  meet  and  combine,  and 
so  set  up  the  equilibrium 

[H+]x[OH-]xft=[H,0]x*',   or 


Now  the  rate  at  which  water  molecules  ionize  is  extremely  slow  as 
compared  with  the  rate  of  their  reunion,  so  that  the  equilibrium  con- 
stant is  numerically  very  small,  and  equilibrium  results  when  almost 
all  the  ions  are  in  the  combined  form.  The  only  ions  left  in  the  solu- 
tion, therefore,  are  K+  and  Cl~,  which  come  to  the  equilibrium  expressed 
in  the  equation  K+  +  Cl~  :p+  KC1 

When  the  solution  is  evaporated,  this  equilibrium  results  in  the  forma- 
tion of  more  and  more  of  the  salt  KC1,  and  it  is  this  which  crystallizes 
from  solution  when  saturation  is  reached.  The  reaction  is  therefore 
practically  completed,  because  one  of  the  possible  equilibrium  products 
is  so  sparingly  ionized  that  it  restores  almost  none  of  its  constituent 
ions  to  the  solution.  But  even  this  small  concentration  of  its  ions, 
H+  and  OH~,  is  of  very  great  importance  under  certain  conditions,  and 
this  will  be  discussed  under  the  topic  Hydrolysis  (see  p.  226). 


202  GENERAL  CHEMISTRY 

The  principle  of  Le  Chatelier.  So  far,  in  discussing  equilibrium,  we 
have  considered  the  factors  which  influence  the  point  of  equilibrium 
and  the  conditions  which  must  prevail  when  equilibrium  is  finally 
reached.  The  question  will  arise,  Supposing  that  when  equilibrium 
has  been  reached,  some  one  factor — say,  pressure  or  temperature  —  is 
changed,  in  which  direction  will  the  equilibrium  shift  ?  In  nearly  all 
cases  the  answer  to  this  question  is  supplied  by  a  very  broad  general- 
ization known  as  the  principle  of  Le  Chatelier.  This  principle  states 
that  when  one  of  the  conditions  attending  equilibrium  is  altered,  that 
one  of  the  opposing  reactions  will  be  aided  which  will  tend  to  restore 
the  original  condition. 

For  example,  let  us  suppose  that  a  gas  is  inclosed  over  water  and 
is  in  equilibrium  with  it  at  a  pressure  of  one  atmosphere.  Let  the 
pressure  be  suddenly  doubled.  An  adjustment  of  equilibrium  will 
follow  which  will  tend  to  restore  the  original  pressure,  and  this  is 
accomplished  by  some  of  the  gas  dissolving  and  so  reducing  the 
pressure.  Again,  let  us  suppose  that  equilibrium  has  been  reached 
in  the  action  of  steam  upon  iron.  If  now  the  partial  pressure  of  the 
steam  is  suddenly  increased,  that  reaction  will  be  aided  which  tends 
to  reduce  it  once  more.  In  other  words,  an  additional  quantity  of 
steam  will  act  upon  the  iron,  thus  reducing  the  steam  pressure.  If 
this  had  been  diminished  instead  of  increased,  an  additional  quantity 
of  hydrogen  would  have  acted  upon  iron  oxide,  to  restore  the  original 
pressure  of  the  steam. 

Suppose  a  hydrated  salt  to  be  in  equilibrium  with  water  vapor  and 
the  anhydrous  salt,  as  in  the  equation 

CaCl2-  2  H20  +=+  CaCl2  +  2  H2O 

and  suppose  heat  is  applied  for  a  few  minutes.  That  reaction  will  be 
promoted  which  absorbs  heat  and  so  reduces  the  temperature  once 
more.  In  other  words,  an  additional  quantity  of  the  hydrate  will  dis- 
sociate. In  an  equilibrium  in  solution,  one  of  the  opposing  reactions  is 
always  endothermic  and  the  other  exothermic.  Heating  such  an  equi- 
librium solution  will  promote  the  endothermic  reaction,  and  cooling 
will  promote  the  exothermic  one.  Since  the  heat  evolved  or  absorbed 
in  the  process  of  ionization  is  usually  very  small,  a  change  of  temper- 
ature does  not  greatly  affect  the  percentage  of  ionization. 


CHAPTER  XVII 

SULFUR;  SELENIUM;  TELLURIUM 

History  and  occurrence.  Sulfur  occupied  a  prominent  place  among 
the  few  elements  known  to  the  ancients,  and  played  an  important  part 
in  the  older  views  concerning  the  composition  of  matter.  It  occurs  in 
nature  in  both  the  free  and  the  combined  condition. 

1.  Free  condition.  In  certain  volcanic  regions,  especially  in  Sicily, 
large  deposits  of  free  sulfur  are  found,  which  until  recent  times  served 
as  the  principal  source  of  the  world's  supply  of  this  element.    Free 
sulfur  also  occurs  in  Japan,  Spain,  Iceland,  Mexico,  and  in  different 
localities  in  the  United  States,  especially  in  Louisiana.    The  deposit 
in  Louisiana  is  remarkable  for  its  magnitude  and  purity.    It  is  located 
about  700  feet  below  the  earth's  surface,  is  circular  in  shape,  about 
one-half  mile  in  diameter,  and  approximately  500  feet  in  thickness. 
Practically  all  of  the  sulfur  now  used  in  the  United  States,  and  some  of 
that  used  in  other  countries  as  well,  is  obtained  from  this  deposit. 

2.  Combined  condition.  Large    quantities    of   sulfur    also    occur   in 
nature  in  the  compounds  known  as  sulfides  and  sulfates.    Some  of 
the  most  important  of  these  are  the  following,  the  names  assigned 
being  the  ones  in  common  use. 


SULFIDES 


Sl'LFATES 


Galena PbS 

Zinc  blende ZnS 

Chalcopyrite CuFeS2 

Pyrites  .  FeS0 


Gypsum CaSO4  •  2  H2O 

Barite BaSO4 

Celestite SrSO4 

Epsom  salt  ....  MgSO4  •  7  H2O 


Since  sulfur  is  a  constituent  of  protein  (p.  307),  it  is  present  quite 
generally  in  vegetable  and  animal  matter,  although  only  in  small 
quantities.  Certain  varieties  of  food,  such  as  the  yolks  of  eggs, 
contain  larger  percentages  of  sulfur. 

Extraction  and  purification  of  sulfur.  The  native  sulfur  found  in 
volcanic  regions  is  mixed  with  more  or  less  rock  and  earthy  material, 
amounting  in  the  case  of  Sicilian  sulfur  to  nearly  75  per  cent  of  the 

203 


204 


GENERAL  CHEMISTRY 


FIG.  79 


entire  mass.  Since  the  element  has  a  relatively  low  melting  point 
(114.8°),  its  separation  from  such  materials  is  accomplished  by  simply 
heating  the  mixture.  The  sulfur  melts  and  drains  away  from  the 
earthy  impurities.  The  crude  sulfur  obtained  in  this  way  is  distilled 
from  an  iron  retort  (Fig.  79,  J),  the  exit  tube  of  which  opens  into  a 
cooling  chamber  (^)  of  brickwork.  When  the  sulfur  vapor  first 
enters  the  cold  chamber,  it  condenses  in  the  form  of  a  powder  known  as 

flowers  of  sulfur.  As  the 
chamber  becomes  warmer 
the  vapor  condenses  to  a 
liquid  and  is  drawn  off 
into  cylindrical  molds,  in 
which  it  solidifies,  form- 


ing the  commercial  prod- 
uct called  brimstone. 

In  Louisiana,  wells  are 
sunk  into  the  deposit  and 
superheated  steam  forced  down  through  suitable  pipes.  The  heat  of 
the  steam  is  sufficient  to  melt  the  sulfur,  which  is  then  forced  through 
tubes  to  the  earth's  surface  by  compressed  air.  A  single  well  often 
produces  daily  as  much  as  500  tons  of  sulfur,  which  is  over  99.5  per 
cent  pure.  About  250,000  tons  are  now  produced  annually  from  this 
deposit,  but  much  larger  quantities  could  be  obtained  if  desired. 

Properties.  Sulfur  is  a  pale  yellow  solid  without  marked  taste  and 
with  but  a  faint  odor.  It  is  insoluble  in  water.  It  melts  when  heated, 
forming  a  thin,  straw-colored  liquid.  If  the  temperature  is  gradually 
raised,  this  liquid  turns  darker  in  color  and  becomes  thicker  until  at 
about  235°  it  is  almost  black,  and  is  so  viscous  that  a  vessel  contain- 
ing it  can  be  inverted  without  danger  of  the  liquid  flowing  out.  At 
higher  temperatures  it  becomes  mobile  again  and  boils  at  444.6°, 
forming  a  yellowish  vapor.  On  cooling,  the  same  changes  take  place 
in  reverse  order. 

Sulfur  exists  in  a  number  of  different  forms,  which  may  be  described 
under  two  general  heads :  namely,  crystalline  sulfur  and  amorphous 
sulfur. 

Crystalline  sulfur.  Sulfur  is  a  dimorphous  element  (p.  82),  forming 
crystals  which  belong  either  to  the  rhombic  or  to  the  monoclinic 
system.  Both  forms  are  insoluble  in  water,  but  are  easily  soluble  in 
the  liquid  known  as  carbon  disulfide. 


SULFUR;   SELENIUM;  TELLURIUM  205 

1.  Rhombic  sulfur.  Sulfur  is  sometimes  found  in  nature  in  crystal- 
line form,  and  these  crystals  always  belong  to  the  rhombic  system.    In 
the  laboratory,  rhombic  crystals  may  be  obtained  by  allowing  a  solu- 
tion of  sulfur  in  carbon  disulfide  to  evaporate  spontaneously.    While 
brimstone  consists,  in  the  main,  of  rhombic  crystals,  one  frequently 
finds,  on  breaking  a  stick  of  it,  that  the  interior  portion  is  made  up  of 
crystals  having  the  form  of  needles,  entirely  different  in  shape  from 
the   rhombic   crystals.    These  needle-shaped  crystals  belong  to  the 
monoclinic  system  and  are  therefore  termed  monoclinic  sulfur. 

2.  Monoclinic  sulfur.  This  form  of  sulfur  is  best  obtained  by  slowly 
cooling  melted  sulfur.    Large  crystals  may  be  obtained  by  melting 
sulfur  in  a  suitable  vessel,  allowing  a  portion  of  the  liquid  to  solidify 
and  then  pouring  off  the  remainder.    The  solid  sulfur  adhering  to  the 
vessel  will  be  found  to  consist  of  distinct  monoclinic  crystals.    These 
crystals  differ  from  the  rhombic  not  only  in  shape  but  in  density  and 
melting  point,  as  shown  in  the  following  table : 


DENSITY 

MELTIXG  POIXT 

Rhombic  sulfur     .... 
Monoclinic  sulfur  .... 

2.06 
1.96 

114.5° 
119.25° 

Relation  of  rhombic  to  monoclinic  sulfur.  Experiments  have  shown 
that  whenever  sulfur  crystallizes  at  ordinary  temperature,  the  rhombic 
form  is  always  obtained.  When  crystallized  at  higher  temperatures,  as 
when  the  sulfur  is  melted  and  allowed  to  cool,  the  monoclinic  form  is 
obtained.  Moreover,  the  temperature  below  which  sulfur  assumes  the 
rhombic  form  and  above  which  it  is  monoclinic  is  a  perfectly  definite 
one,  namely,  95.5°.  At  this  transition  point  (p.  81)  the  two  forms  of 
crystals  remain  unchanged  when  in  contact  with  each  other.  If  heated 
above  95.5°,  the  rhombic  form  gradually  changes  into  the  monoclinic 
form;  if  cooled  below  95.5°,  the  monoclinic  gradually  changes  into 
the  rhombic  form.  This  change  of  one  form  into  the  other  ordinarily 
takes  place  very  slowly,  so  that  some  days  may  be  required  before  the 
change  is  complete.  From  these  facts  one  can  readily  understand  why 
the  crystals  found  in  nature  always  have  the  rhombic  form. 

Amorphous  sulfur.  In  discussing  the  physical  properties  of  sul- 
fur, attention  was  called  to  the  fact  that,  when  heated,  sulfur  melts, 
forming  a  pale  yellow,  mobile  liquid,  which  at  a  higher  temperature 
becomes  dark  and  viscous.  The  American  chemist  Alexander  Smith 


206 


GENERAL  CHEMISTKY 


has  shown  that  at  intermediate  temperatures  the  liquid  obtained  con- 
sists of  varying  amounts  of  the  mobile  and  the  viscous  liquid  in  equilib- 
rium with  each  other.  To  distinguish  the  mobile  liquid  from  the  thick, 
viscous  liquid,  Smith  has  proposed  that  the  former  be  represented  by 
the  symbol  87  and  the  latter  by  S/z.  If  the  molten  sulfur  is  heated 
to  boiling  and  poured  into  cold  water,  the  sudden  chilling  prevents 
the  crystallization  of  the  viscous  liquid,  so  that  an  amorphous,  dough- 
like  product  is  obtained.  This  form  is  insoluble  in  carbon  disulfide, 
and  is  known  as  plastic  sulfur. 

The  formation  of  plastic  sulfur  is  shown  in  a  very  striking  manner  by 
distilling  sulfur  from  a  small,  short-necked  retort  (Fig.  80)  and  allowing 

the  distillate  to  run  into  cold  water. 
On  standing  at  ordinary  tempera- 
tures, plastic  sulfur  changes  in  part 
into  rhombic  crystals.  The  crystal- 
line portion  may  then  be  dissolved 
in  carbon  disulfide,  leaving  an  insol- 
uble residue  which  is  amorphous  and 
apparently  permanent.  The  same 
amorphous  form  is  present  in  flowers 
of  sulfur,  and  is  left  as  an  insoluble 
residue  when  the  latter  is  treated 
with  carbon  disulfide. 
Milk  of  sulfur  (sulfur  lac).  When  an  acid  is  added  to  aqueous 
solutions  of  certain  compounds  rich  in  sulfur,  a  portion  of  the  element 
is  set  free  and,  being  insoluble,  separates  in  the  form  of  a  white  pow- 
der which  imparts  to  the  entire  liquid  a  milky  appearance ;  hence  the 
name  milk  of  sulfur.  This  form  consists  chiefly  of  rhombic  crystals. 
Like  all  substances  formed  by  precipitation,  it  is  in  a  fine  state  of 
division,  and,  because  of  this  fact,  more  readily  undergoes  chemical 
changes  (p.  28).  It  is  the  form  commonly  used  in  medicine. 

Chemical  conduct.  Sulfur  combines  directly  with  many  of  the  ele- 
ments, metals  as  well  as  nonmetals.  The  compounds  formed  by  its 
union  with  -some  other  element  are  termed  sulfides. 

1.  Action  upon  metals.  Most  metals,  when  heated  with  sulfur,  com- 
bine directly  with  it,  forming  metallic  sulfides.  In  some  cases  the 
action  is  so  energetic  that  the  mass  becomes  incandescent,  as  has 
been  seen  in  the  case  of  the  reaction  of  iron  with  sulfur  (p.  8).  This 
property  recalls  the  action  of  oxygen  upon  metals,  and  in  general  it 


FIG.  80 


SULFUR;   SELENIUM;   TELLURIUM  207 

may  be  said  that  the  metals  which  combine  readily  with  oxygen  are 
apt  to  combine  quite  readily  with  sulfur. 

2.  Action  upon  nonmetals.  Under  suitable  conditions  sulfur  combines 
with  most  of  the  nonmetals.    Thus,  with  hydrogen  it  forms  the  gas 
hydrogen  sulfide  (H2S) ;  with  carbon  it  forms  carbon  disulfide  (CS2), 
a  heavy,  colorless  liquid ;  with  chlorine  it  forms  the  liquid  chlorides 
S2C12  and  SC14.    At  ordinary  temperatures  and  in  the  presence  of 
moisture,  sulfur  reacts  slowly  with  oxygen,  forming  sulf uric  acid ;  at 
higher  temperatures  it  burns,  forming  sulfur  dioxide  (SO2),  a  gaseous 
compound  having  the  well-known  odor  of  burning  sulfur. 

3.  Action  with  oxidizing  agents.    When  acted  upon  by  oxidizing 
agents  in  the  presence  of  water,  sulfur  is  converted  into  sulfuric  acid. 
Since  very  accurate  methods  are  known  for  the  determination  of  sul- 
furic acid,  this  reaction  is  the  one  generally  used  hi  determining  the 
percentage  of  sulfur  present  in  any  substance. 

Uses  of  sulfur.  Large  quantities  of  sulfur  are  used  in  the  manu- 
facture of  its  compounds,  such  as  sulfuric  acid  and  carbon  disulfide. 
It  is  also  used  extensively  in  the  manufacture  of  gunpowder,  matches, 
vulcanized  rubber,  and  insecticides. 

COMPOUNDS  OF  SULFUR  WITH  HYDROGEN 

The  following  compounds  of  sulfur  and  hydrogen  are  known: 
hydrogen  sulfide  (H2S),  a  foul-smelling  gas;  hydrogen  persulfide,  a 
liquid  which  is  probably  a  mixture  of  the  sulfides  H2S4  and  H2S6. 

Hydrogen  sulfide  (H2S).  Hydrogen  sulfide  is  present  in  the  vapors 
issuing  from  volcanoes.  Dissolved  in  water,  it  constitutes  the  so-called 
sulfur  waters  of  common  occurrence.  It  is  formed  when  organic  matter 
containing  sulfur  undergoes  decay,  and  the  disagreeable  odor  attend- 
ing such  changes  is  often  partly  due  to  the  presence  of  this  gas. 

Preparation.  Hydrogen  sulfide  may  be  prepared  by  the  two  follow- 
ing general  methods : 

1.  Synthetic  method.  When  hydrogen  is  bubbled  through  boiling 
sulfur,  a  portion  of  the  gas  combines  directly  with  the  sulfur  to  form 
hydrogen  sulfide.  It  is  likewise  formed  when  hydrogen  is  passed  over 
certain  sulfides  heated  to  a  definite  temperature.  It  may  be  added 
that  both  these  reactions  are  reversible,  as  shown  in  the  following 
equations:  H2  +  S  ^±  H2S 

4  H2  +  Fe8S4  +±.  3  Fe  +  4  H2S 


208 


GENERAL  CHEMISTRY 


2.  Laboratory  method.  Hydrogen  sulfide  is  most  conveniently  pre- 
pared in  the  laboratory  by  the  action  of  either  dilute  sulfuric  or 
hydrochloric  acid  upon  the  sulfide  of  a  metal.  Iron  sulfide  (FeS)  is 
usually  employed,  although  the  resulting  hydrogen  sulfide  is  not 
entirely  pure.  If  a  greater  purity  is  desired,  one  may  use  either  sul- 
fide of  antimony  (Sb.2S3),  which  is  found  in  nature  in  a  state  of  great 
purity,  or  sodium  sulfide,  which  can  be  prepared  in  the  laboratory  in 
a  pure  state.  When  the  acid  is  added  to  the  sulfide,  the  hydrogen 
ions  derived  from  the  acid  and  the  sulfur  ions  derived  from  the  sul- 
fide tend  to  form  an  equilibrium  with  the  hydrogen  sulfide  resulting 

from  their  union  :         TT+      TT+      Q__     _    u  c 
H^  +  H+  +  b       <     >  tip 

Before  equilibrium  is  reached,  the  water  present  becomes  saturated 
with  the  gas,  which  then  escapes  as  fast  as  formed,  so  that  the  re- 
action continues  until  either  the  acid  or  the  sulfide  is  practically 
all  consumed.  The  following  equations  serve  to  show  the  proportion 
in  which  the  sulfide  and  acids  interact,  as  well  as  the  final  products 
formed  in  each  case  : 


+  g 


+ 


FIG.  81 


To  prepare  hydrogen  sulfide  in  the  laboratory, 
iron  sulfide  is  placed  in  a  flask  A  (Fig.  81)  and 
dilute  acid  added  drop  by  drop  from  the  separa- 
tory  funnel  B.  The  hydrogen  sulfide  formed 
escapes  through  the  tube  C  and  is  collected  in 
cylinders,  as  shown  in  the  figure.  The  Kipp 
generator  (p.  38)  serves  as  a  more  convenient 
form  of  apparatus  for  generating  the  gas.  By  the 
evaporation  of  the  liquid  left  in  the  generating 
flask  there  will  be  obtained  either  chloride  of  iron 
(FeCl2)  or  sulfate  of  iron  (FeSO4),  according  to 
the  acid  used. 


Properties.  Hydrogen  sulfide  is  a  colorless  gas  having  a  mild,  dis- 
agreeable taste  and  an  offensive  odor.  It  is  1.18  times  as  heavy  as 
air.  The  gas  may  be  readily  condensed  to  a  colorless  liquid  which 
boils  at  -  61.8°  and  solidifies  at  -  86°.  One  volume  of  water  at  15° 
dissolves  3.05  volumes  of  the  gas.  When  this  solution  is  heated  to 
boiling,  the  gas  is  all  expelled.  In  pure  form  it  acts  as  a  violent 
poison  and,  even  when  diluted  largely  with  air,  produces  headache, 
dizziness,  and  nausea.  Fortunately  its  extremely  disagreeable  odor 
gives  warning  of  its  presence. 


SULFUR;   SELENIUM;  TELLURIUM  209 

Chemical  conduct.  1.  Acid  properties.  In  aqueous  solution,  hydrogen 
sulfide  is  slightly  dissociated,  giving  hydrogen  ions.  The  solution 
therefore  acts  as  a  weak  acid  and  is  known  as  hydrosulfuric  acid.  It 
possesses  the  general  properties  of  an  acid,  turning  blue  litmus  red 
and  neutralizing  bases  with  the  formation  of  sulfides. 

2.  Action  of  heat.  When  heated  to  a  high  temperature,  hydrogen 
sulfide  is  decomposed  into  its  elements,  the  speed  of  decomposition 
being  marked  at  500°. 

3.  Action  of  oxygen.    When  a  solution  of  hydrogen  sulfide  in  water 
(hydrosulfuric  acid)  is  exposed  to  the  air,  the  hydrogen  of  the  sulfide 
unites  with  oxygen  to  form  water,  while  the  sulfur  is  liberated  and 
settles  to  the  bottom  of  the  liquid.   In  this  way  are  formed  the  deposits 
of  the  element  found  about  sulfur  springs.    At  a  high  temperature; 
hydrogen  sulfide  burns  readily  in  either  oxygen  or  air,  according  to 
the  equation  2  H2S  +  3  O2  =  2  H2O  +  2  SO2 

When  there  is  not  sufficient  oxygen  to  combine  with  both  the  sulfur 
and  the  hydrogen,  the  latter  element  combines  with  the  oxygen  and  the 
sulfur  is  set  free:  2H,S  +  O2=  2H2O  +  2S 

4.  Reducing  action.    Because  of  the  hydrogen  present,  together  with 
the  ease  with  which  it  is  given  up  in  contact  with  an  oxidizing  agent, 
hydrogen  sulfide  acts  as  a  strong  reducing  agent.    Thus,  when  it  is 
bubbled  through  concentrated  nitric  or  sulfuric  acid,  both  of  which 
are  strong  oxidizing  agents,  the  hydrogen  of  the  sulfide  combines  with 
a  portion  of  the  oxygen  of  the  acid  to  form  water,  the  acid  being  at 
the  same  time  reduced. 

A  common  method  of  drying  gases  consists  in  bubbling  them  through  concen- 
trated sulfuric  acid,  which  absorbs  the  moisture.  It  is  evident,  however,  from  the 
statements  just  made,  that  this  method  cannot  be  used  for  drying  hydrogen  sulfide. 

5.  Action  upon  metals.    The  action  of  hydrogen  sulfide  upon  a  num- 
ber of  the  metals  is  very  similar  to  that  of  water  under  like  conditions, 
resulting  in  the  displacement  of  its  hydrogen  by  the  metal.    Thus, 
when  it  is  passed  over  iron  filings  heated  in  a  tube,  the  reaction  repre- 
sented by  the  following  equation  takes  place : 

3Fe  +  4H2S >.Fe8S4  +  4H2 

Under  similar  circumstances  steam  acts  according  to  the  following 
equation:  8  Fe  +  4  H.O  — vFe.O.  +  4  H. 


210 


GENERAL  CHEMISTRY 


Salts  of  hydrosulfuric  acid ;  sulfides.  The  salts  of  hydrosulfuric 
acid,  or  sulfides,  form  an  important  class  of  compounds.  They  are  all 
solids ;  most  of  them  are  insoluble  in  water,  while  some  are  insoluble 
even  in  acids.  As  prepared  in  the  laboratory,  some  of  these  salts,  such 
as  copper  sulfide  (CuS)  and  silver  sulfide  (Ag2S),  are  black  ;  others,  as 
cadmium  sulfide  (CdS)  and  arsenic  sulfide  (As2S3),  are  yellow ;  while 
zinc  sulfide  (ZnS)  is  white.  Many  of  these  sulfides  are  found  in  nature 
(p.  203),  some  of  them  constituting  the  compounds  from  which  the 
metals  are  obtained  on  a  commercial  scale.  They  will  be  frequently 
mentioned  in  connection  with  the  metals. 

Preparation  of  the  sulfides.  The  soluble  sulfides,  like  those  of  sodium 
and  potassium,  are  most  readily  prepared  by  treating  the  respective 
hydroxides  of  these  metals  with  hydrosulfuric  acid.  Both  the  acid  and 
normal  salts  may  be  obtained  in  this  way: 

NaOH  4-  H2S  =  NaHS  +  H2O 
2  NaOH  +  H2S  =  Na2S  +  H2O 

The  insoluble  sulfides  may  be  prepared  by  heating  the  metals  with 
sulfur,  although  the  general  and  more  convenient  method  of  preparing 
them  consists  in  passing  hydrogen  sulfide  into  aqueous  solutions  of 
metallic  compounds.  Thus,  copper  sulfide  may  be  easily  prepared  by 
dissolving  copper  sulfate  (CuSO4)  in  water  and  passing  hydrogen  sul- 
fide into  the  solution.  The  copper  sulfide,  being  insoluble,  precipitates 
as  fast  as  formed,  and  may  be  removed  from  the  liquid  by  filtration. 
The  reaction  which  takes  place  is  expressed  in  the  following  equation : 

H2S  +  CuSO4  =±=  CuS  +  H2SO4 

The  preparation 
of  these  sulfides  as 
carried  out  in  the 
laboratory  may  be 
illustrated  in  the 
following  way : 

Hydrogen  sulfide  is 
generated  in  a  Kipp 
apparatus  A  (Fig.  82) 
and  is  passed  succes- 
sively into  bottles  B, 

1 IG.  o2  ^.     _      _.  . 

C,  D,  E,  containing, 

respectively,  the  aqueous  solutions  of  silver  nitrate,  cadmium  sulfate,  zinc  acetate, 
and  sodium  hydroxide.  As  the  gas  bubbles  through  the  solutions  there  is  formed 


SULFUR;   SELENIUM;   TELLURIUM  211 

black  silver  sulfide  (Ag2S)  in  B,  yellow  cadmium  sulfide  (CdS)  in  C,  white  zinc 
sulfide  (ZnS)  in  D.  No  precipitate  is  produced  in  E,  for  although  sodium  sulfide 
is'  formed,  it  is  soluble  in  water  and  therefore  does  not  separate. 

The  persulfides.  In  addition  to  the  ordinary  sulfides,  which  are  salts 
of  hydrosulfuric  acid,  some  of  the  metals  form  persulfides.  As  is  indi- 
cated by  the  name,  these  compounds  contain  a  higher  percentage  of  sul- 
fur than  the  ordinary  sulfides.  The  persulfides  of  sodium,  potassium, 
and  calcium  are  formed,  together  with  other  products,  when  aqueous 
solutions  of  the  hydroxides  of  these  metals  are  heated  with  sulfur. 
The  exact  composition  of  these  persulfides  is  not  known  with  cer- 
tainty. They  are  all  unstable,  readily  decomposing  into  the  ordinary 
sulfides  and  free  sulfur.  Because  of  this  property  some  of  the  per- 
sulfides are  coming  into  general  use  as  insecticides,  the  free  sulfur 
being  effective  for  this  purpose.  Thus,  the  lime-sulfur  spray  that  is 
being  used  so  extensively  at  present  for  destroying  insects  injurious 
to  trees  is  simply  the  red  solution  of  calcium  persulfide  (probably  a 
mixture  of  CaS4  and  CaS5),  formed  by  heating  calcium  hydroxide 
with  sulfur  and  water,  as  described  above. 

When  an  acid  such  as  hydrochloric  is  added  to  a  solution  of  the 
persulfide,  a  portion  of  the  sulfur  present  is  evolved  as  hydrogen  sulfide, 
and  a  portion  is  liberated  in  the  free  state  (milk  of  sulfur,  p.  206). 
If,  however,  in  place  of  adding  the  acid  to  the  solution  of  the  per- 
sulfide, the  latter  is  added  slowly  to  the  acid,  a  heavy  yellow  oil 
separates.  This  oil  is  termed  hydrogen  persulfide  and  is  probably 
a  mixture  of  the  compounds  H2S4  and  H2Sg.  This  mixture  is  very 
unstable,  rapidly  decomposing  into  hydrogen  sulfide  and  free  sulfur. 


OXIDES  AND  OXYGEN  ACIDS  OF  SULFUR 
Sulfur  forms  the  following  compounds  with  oxygen : 

Sulfur  dioxide  (SO2)  :  a  colorless  gas. 
Sulfur  trioxide  (SO3)  :  a  colorless  liquid  boiling  at  46°. 
Sulfur  sesquioxide  (S2O3)  :  a  bluish-green,  crystalline  solid. 
Sulfur  hexoxide  (S2O6)  :  a  white  solid  melting  at  50°. 
Sulfur  heptoxide  (S2O?)  :  a  viscous  liquid. 

Of  these  oxides  the  first  two  are  by  far  the  most  important  and  best 
known.  They  are  both  acid  anhydrides,  combining  with  water  to  form 
sulfurous  acid  and  sulfuric  acid  respectively. 


212  GENERAL  CHEMISTRY 

Sulfur  dioxide  (sulfurous  anhydride)  (S02).  This  is  the  well-known 
gas  resulting  from  the  combustion  of  sulfur.  It  was  first  obtained  in 
the  pure  state  and  recognized  as.  a  definite  compound  by  Priestley 
in  1775.  It  occurs  in  nature  in  the  gas  issuing  from  volcanoes  and 
in  solution  in  the  waters  of  some  springs. 

Preparation.  Sulfur  dioxide  is  prepared  by  three  general  methods : 

1.  By  the  combustion  of  sulfur  or  a  metallic  sulfide.    In  either  case 
the  sulfur  is  converted  into  sulfur  dioxide : 

S  +  O2  =  SO2 
2  ZnS  +  3  O2  =  2  ZnO  +  2  SO2 

The  enormous  quantities  of  sulfur  dioxide  used  in  the  manufacture 
of  sulfuric  acid  are  prepared  by  this  general  method. 

2.  By  the  reduction  of  sulfuric  acid.    When  concentrated  sulfuric  acid 
is  heated  with  certain  metals,  such  as  copper,  a  part  of  the  acid  is 
reduced  to  sulfurous  acid.   The  latter  compound  then  decomposes  into 
sulfur  dioxide  and  water,  the  complete  equation  being  as  follows : 

Cu  +  2  H2S04  =  CuS04  +  S02  +  2  H2O 

A  similar  reaction  takes  place  when  carbon  is  heated  with  sulfuric  acid : 
C  +  2  H2S04  =  C02  +  2  S02  +  2  H2O 

3.  By  the  action  of  acids  upon  a  sulfite.    Sulfites  are  salts  of  sulfurous 
acid  (H2SO3).    When  an  acid,  such  as  hydrochloric  acid,  is  added  to  a 
sulfite,  sulfurous  acid  is  formed,  which  decomposes  into  water  and  sul- 
fur dioxide.   The  reactions  are  expressed  in  the  following  equations  : 

Na2S03  -f  2  HC1  +=*  2  NaCl  +  H2SO3  (1) 

Explanation  of  the  reaction.  In  the  action  of  hydrochloric  acid  upon  sodium 
sulfite,  as  expressed  in  these  equations,  we  have  two  reversible  reactions  depend- 
ing upon  each  other.  It  might  be  expected  that  the  reaction  expressed  in  equa- 
tion (1)  would  result  in  an  equilibrium,  since  none  of  the  substances  represented 
in  the  equation  are  insoluble  or  volatile  in  the  presence  of  water.  The  sulfurous 
acid,  however,  decomposes  as  fast  as  it  forms,  according  to  equation  (2),  the  result- 
ing sulfur  dioxide  escaping  in  the  form  of  a  gas.  The  reaction  continues,  there- 
fore, until  practically  all  the  sodium  sulfite  has  been  decomposed.  Since  sulfur 
dioxide  is  quite  soluble  in  water,  it  is  evident  that  the  reaction  should  be  carried 
out  in  the  presence  of  as  little  water  as  possible ;  otherwise  a  proportionately 
larger  quantity  of  sulfur  dioxide  will  remain  in  solution,  and  the  reaction  will  not 
reach  the  same  degree  of  completion. 


SULFUR;   SELENIUM;   TELLURIUM 


213 


FIG.  83 


Properties.  Sulfur  dioxide  is  a  colorless  gas  and  has  the  peculiar 
irritating  odor  so  noticeable  when  sulfur  is  burned.  It  is  a  heavy  gas, 
being  2.2  times  as  heavy  as  air.  Under  standard  conditions  1  volume 
of  water  dissolves  about  80  volumes  of  the  gas.  It  is  easily  liquefied, 
a  freezing  mixture  of  ice  and  salt  being  sufficient  to  effect  condensa- 
tion under  atmospheric  pressure.  The  resulting  liquid  is  colorless,  boils 
at  —10.1°,  and  freezes  to  a 
snowlike  solid  mass  which 
melts  at  —  76°.  Liquid  sul- 
fur dioxide  is  a  commercial 
product,  being  stored  in 
strong  glass  siphon  bottles 
or  in  metal  cylinders. 

The  condensation  of  the 
sulfur  dioxide  to  a  liquid  can 
be  accomplished  in  the  follow- 
ing way :  Pure  sulfur  dioxide 
is  generated  in  A  (Fig.  83)  and 
conducted  into  an  empty  flask 
B  surrounded  by  ice  and  salt 

in  the  ratio  3  : 1  by  weight.  The  low  temperature  produced  by  this  mixture  is 
sufficient  to  condense  a  portion  of  the  gas  to  a  liquid,  which  collects  in  the  bottle  B. 

Chemical  conduct.  Sulfur  dioxide  combines  with  a  number  of  other 
substances.  Thus,  with  chlorine  it  forms  the  compound  known  as 
sulfuryl  chloride  (SO2C12),  while  with  oxygen  it  forms  sulfur  trioxide. 
It  also  reacts  with  hydrogen  sulfide  to  form  water  and  free  sulfur : 

2H2S  +  S02=2H20  +  3S 

Since  both  hydrogen  sulfide  and  sulfur  dioxide  are  present  in  the 
gases  issuing  from  volcanoes,  it  is  probable  that  the  large  deposits  of 
sulfur  occurring  in  volcanic  regions  have  resulted  from  the  interaction 
of  these  two  gases,  according  to  the  above  equation.  From  these  state- 
ments it  will  be  seen  that  under  some  conditions  the  gas  takes  up 
oxygen  to  form  a  higher  oxide,  acting  as  a  reducing  agent ;  under 
other  conditions  it  gives  up  oxygen,  acting  as  an  oxidizing  agent.  It 
will  be  remembered  that  nitric  oxide  acts  in  a  similar  way. 

A  characteristic  property  of  sulfur  dioxide  is  its  conduct  toward 
water,  with  which  it  unites  to  form  sulfurous  acid. 

Sulfurous  acid  (H2S03).  When  sulfur  dioxide  is  passed  into  water, 
some  of  the  gas  combines  with  water  to  form  sulfurous  acid  (H2SOg), 


214  GENERAL  CHEMISTRY 

while  the  remainder  is  held  in  a  state  of  solution.  The  sulfurous  acid 
formed  is  in  equilibrium,  on  the  one  hand  with  water  and  dissolved 
sulfur  dioxide,  and  on  the  other  hand  with  the  ions  H+  and  HSO8~, 
resulting  from  the  ionization  of  a  portion  of  the  acid  : 

H20  +  S02  j=±  H2S08  +=±  H+,  HS08- 

When  heated,  this  liquid  acts  as  if  it  were  simply  a  solution  of  sul- 
fur dioxide  in  water,  all  the  sulfur  being  evolved  as  sulfur  dioxide. 
Toward  a  base,  on  the  other  hand,  it  acts  simply  as  a  solution  of 
sulfurous  acid  (compare  with  aqua  ammonia,  p.  172). 

Because  of  its  unstable  character,  sulfurous  acid  can  be  obtained 
only  in  the  form  of  a  dilute  solution.  This  solution  has  the  following 
properties  : 

1.  Acid  properties.  In  aqueous  solutions  the  compound  forms  the 
ions  H+  and  HSO8~.    In  very  dilute  solutions  some  of  the  latter  ions 
may  decompose  further  into  the  ions  H+  and  SO8~~.    The  solution 
has  all  the  properties  of  a  weak  acid. 

2.  Reducing  properties.    Sulfurous  acid  is  a  good  reducing  agent, 
taking  up   oxygen  either  from  the  air  or  from  substances  rich  in 
oxygen,  and  changing  into  sulfuric  acid,  as  shown  in  the  following 
equations  :  2  H2SO8  +  O2  =  2  H2SO4 


In   this   respect    it   resembles   its  anhydride   SO2,   which   is  also    a 
reducing  agent. 

Because  of  this  tendency  of  sulfurous  acid  to  combine  with  oxygen,  a  solution 
of  the  acid,  unless  freshly  prepared,  always  contains  more  or  less  sulfuric  acid. 
It  has  been  found  that  the  speed  of  the  reaction  by  which  sulfurous  acid  is 
changed  into  sulfuric  acid  is  greatly  diminished  in  the  presence  of  a  trace  of 
sugar  or  glycerin.  These  latter  substances  act  simply  as  catalytic  agents.  While 
catalyzers  are  generally  employed  to  increase  the  speed  of  a  reaction,  it  will  be 
noted  that  in  this  case  the  sugar  and  glycerin  have  just  the  opposite  effect, 
namely,  to  diminish  the  speed  of  the  reaction  (see  p.  192). 

3.  Bleaching  properties.  Sulfurous  acid  acts  upon  many  organic  dyes, 
changing  them  into  colorless  compounds.  It  is  therefore  useful  as  a 
bleaching  agent,  and  is  used  especially  to  bleach  such  substances  as 
paper  and  straw  goods,  the  texture  of  which  would  be  injured  by 
a  more  powerful  bleaching  agent,  like  chlorine. 

The  bleaching  properties  of  sulfurous  acid  may  be  shown  by  bringing  a  small 
dish  of  burning  sulfur  under  a  bell  jar  (Fig.  84)  in  which  has  been  placed  some 
highly  colored  flower  thoroughly  moistened  with  water.  The  sulfur  dioxide 


SULFUK;  SELENIUM;   TELLURIUM  215 

combines  with  the  moisture  to  form  sulfurous  acid,  which  slowly  bleaches  the 
flower.  The  reactions  involved  in  these  changes  are  not  thoroughly  understood. 
In  some  cases  the  sulfurous  acid  apparently  combines  directly  with  the  coloring 
matter  to  form  more  or  less  unstable  colorless  com- 
pounds. The  gradual  decomposition  of  these  compounds 
into  their  original  constituents  would  account  for  the 
fact  that  some  substances  bleached  by  sulfurous  acid 
gradually  regain  their  original  color.  In  other  cases 
the  bleaching  properties  of  the  acid  seem  to  be  due  to 
its  reducing  action,  whereby  oxygen  is  removed  from  the 
water  present,  while  the  hydrogen  liberated  combines 
with  the  coloring  matter  to  form  colorless  compounds. 

4.   Antiseptic  and  preservative  properties.    Sul- 
furous acid  destroys  many  microorganisms  and 

may  therefore  be  used  as  a  disinfectant  and  also  as  a  preservative,  to 
prevent  such  changes  as  decay  and  fermentation,  which  are  caused  by 
these  organisms. 

Structural  formula  of  sulfurous  acid.  Both  of  the  following  struc- 
tural formulas  have  been  assigned  to  the  acid,  but  its  chemical  reac- 
tions are  not  such  as  to  show  which  of  the  two  is  correct. 

H0>s  H0>s^0 

H0>  H>S*0 

Uses  of  sulfurous  acid.  Sulfurous  acid  is  used  mainly  in  the  prepa- 
ration of  sulfites  and  as  a  bleaching  agent  and  preservative.  Formerly 
it  was  often  used  for  disinfecting  rooms  after  the  occurrence  of 
contagious  diseases,  but  it  is  now  largely  replaced  by  formaldehyde 
(p.  304),  which  is  more  effective  for  the  destruction  of  micro- 
organisms. It  is  used  to  a  considerable  extent  in  certain  foods  such 
as  canned  corn,  dried  fruits,  sirups,  and  wines,  serving  not  only  as  a 
preservative  but  also  as  a  bleaching  agent,  to  remove  objectionable 
colors.  Whether  or  not  its  use  in  foods  should  be  permitted  is  a 
much-debated  question. 

Salts  of  sulfurous  acid ;  sulfites.  Being  dibasic,  sulfurous  acid  forms 
both  acid  and  normal  salts.  Thus,  with  sodium  it  forms  the  salts 
NaHSO8  and  Na2SO3.  The  sulfites  of  most  of  the  metals  are  known, 
although  some  of  them  have  not  been  obtained  in  pure  condition. 
With  the  exception  of  the  sulfites  of  sodium  and  potassium,  the  nor- 
mal sulfites  are  either  insoluble  in  water  or  nearly  so,  while  the  acid 
sulfites  are  soluble.  The  normal  salts  can  be  prepared  either  by  the 
action  of  sulfurous  acid  upon  the  hydroxides  of  the  metals  or  by 


216 


GENERAL   CHEMISTRY 


the  general  precipitation  method  used  for  preparing  insoluble  com- 
pounds. The  sulfites  are  all  solid  substances  and,  like  sulfurous 
acid  itself,  combine  readily  with  oxygen,  forming  the  corresponding 
sulfates.  They  are  therefore  good  reducing  agents.  Because  of  this 
property,  unless  freshly  prepared,  they  are  apt  to  contain  more  or 
less  of  the  corresponding  sulfates.  Calcium  acid  sulfite  is  largely  used 
in  the  manufacture  of  paper  from  wood,  since  it  dissolves .  the  objec- 
tionable constituent  (lignin)  of  the  wood,  leaving  the  pure  cellulose, 
which  is  the  material  desired  for  the  manufacture  of  paper. 

Sulfur  trioxide  (sulfuric  anhydride)  (S03).  When  sulfur  is  burned  in 
oxygen,  minute  quantities  of  sulfur  trioxide  are  formed  along  with 
the  sulfur  dioxide.  Likewise,  when  sulfur  dioxide  and  oxygen  are 
heated  together,  combination  takes  place,  but  the  speed  of  the  reaction 
is  so  slow  that  only  traces  of  the  trioxide  result.  In  the  presence  of 
a  catalytic  agent,  however,  such  as  finely  divided  platinum,  the  speed 
is  greatly  increased,  and  in  this  way  sulfur  trioxide  can  be  obtained 
in  quantities.  The  reaction  is  a  reversible  one,  as  is  shown  in  the 
following  equation : 


2SOi+'P1 


2  SO, 


The  largest  yield  of  sulfur  trioxide  is  obtained  when  the  reaction  is 
carried  out  at  approximately  400°,  at  which  temperature  about  98 
per  cent  of  the  sulfur  dioxide  combines  with  oxygen. 

The  preparation  of  the  trioxide  by  this  method  can  be  carried  out  in  the 
laboratory  as  follows :  The  platinum  used  as  a  catalytic  agent  is  prepared  by 
moistening  asbestos  fiber  in  a  solution  of  chloroplatinic  acid  and  igniting  it  in  a 


FIG.  85 


flame,  whereby  the  platinum  compound  is  reduced  to  metallic  platinum.  The 
fiber  containing  the  finely  divided  platinum  is  placed  in  a  tube  of  hard  glass  A 
(Fig.  85),  which  is  then  heated  to  about  400°,  while  equal  volumes  of  sulfur 
dioxide  and  oxygen,  previously  dried  by  bubbling  them  through  sulfuric  acid 


SULFUR;   SELENIUM;  TELLURIUM  217 

(contained  in  bottles  B  and  C),  are  passed  into  the  tube.  As  this  mixture  comes 
in  contact  with  the  catalytic  agent,  combination  takes  place,  and  the  resulting 
sulfur  trioxide  escapes  from  the  jet  at  the  end  of  the  tube  and  may  be  condensed 
by  surrounding  the  receiving  tube  D  with  a  freezing  mixture. 

Properties.  Sulfur  trioxide  is  a  colorless  liquid  which  solidifies  at 
about  15°  and  boils  at  46°.  It  readily  polymerizes  (p.  187),  espe- 
cially in  the  presence  of  a  trace  of  moisture,  forming  sulfur  hexoxide 
(S2Og),  a  white,  crystalline  mass  somewhat  resembling  asbestos  in 
appearance.  Toward  a  substance  having  a  strong  affinity  for  oxygen, 
sulfur  trioxide  acts  as  an  oxidizing  agent,  giving  up  one  third  of  its 
oxygen  and  being  reduced  to  sulfur  dioxide.  It  is  characterized  espe- 
cially by  its  strong  affinity  for  water,  with  which  it  combines  to  form 
sulfuric  acid. 

Sulfuric  acid  (H2SOJ.  Sulfuric  acid  has  long  been  known,  and  was 
one  of  the  most  important  reagents  employed  by  the  alchemists.  It  is 
by  far  the  most  largely  used  of  all  the  acids.  Not  only  is  it  one  of  the 
most  common  reagents  in  the  laboratory,  but  enormous  quantities  of 
it  are  consumed  in  the  industries,  especially  in  the  manufacture  of 
fertilizers,  the  refining  of  petroleum,  and  in  cleaning  scale  from  iron 
and  steel. 

Manufacture  of  sulfuric  acid.  Two  general  methods  for  the  manu- 
facture of  sulfuric  acid  are  in  use  at  the  present  time.  These  are 
known  as  the  contact  process  and  the  lead-chamber  process. 

1.  Contact  process.  The  reactions  taking  place  in  this  process  are 
represented  by  the  following  equations: 

S  +  02=S02  (1) 

2S02+02=2S08  (2) 

S08+H20  =  H2SO,  (3) 

Sulfur  dioxide  is  prepared  according  to  equation  (1),  by  burning  sul- 
fur or  some  sulfide  such  as  iron  pyrites  (FeS2),  in  air.  The  result- 
ing sulfur  dioxide,  together  with  sufficient  air  to  furnish  the  necessary 
oxygen,  is  conducted  through  iron  tubes  filled  with  some  porous 
material  (asbestos  or  sodium  sulfate),  through  which  a  suitable  cata- 
lytic agent,  such  as  platinum  or  iron  oxide,  is  interspersed,  the  material 
being  kept  at  about  400°.  Under  these  conditions  sulfur  trioxide 
is  formed  according  to  equation  (2).  The  resulting,  sulfur  trioxide 
is  then  brought  into  contact  with  water,  with  which  it  unites  to  form 
sulfuric  acid  according  to  equation  (3). 


218  GENERAL  CHEMISTRY 

The  only  part  of  the  process  which  is  difficult  to  carry  out  on  a  commercial 
scale  is  the  formation  of  the  sulfur  trioxide.  It  has  long  been  known  that  sul- 
fur dioxide  and  oxygen  combine  when  passed  over  finely  divided  platinum,  but 
the  cost  of  platinum,  together  with  the  poor  yield  of  sulfur  trioxide  obtained, 
made  the  process  an  impracticable  one.  A  study  of  the  conditions  under  which 
the  reaction  takes  place  resulted  in  improvements  in  the  process,  until  finally,  in 
1901,  the  German  chemist  Knietsch  succeeded  in  overcoming  the  difficulties  to 
such  an  extent  as  to  make  the  process  a  commercial  success  for  the  manufacture 
of  the  pure,  concentrated  acid.  While  platinum  is  the  most  effective  catalytic 
agent  for  the  process,  it  is  very  expensive,  its  commercial  value  at  the  present 
time  being  greater  than  that  of  gold.  This  has  led  to  the  use  of  other  catalytic 
agents,  among  which  iron  oxide  appears  to  be  the  best. 

Some  of  the  more  important  conditions  necessary  for  making  the  process  a 
success  are  the  following :  (1)  The  sulfur  dioxide  must  be  free  from  dust  parti- 
cles and  other  impurities,  such  as  the  traces  of  oxides  of  arsenic,  which  are 
likely  to  be  formed  in  the  combustion  of  iron  pyrites  or  other  sulfides,  which 
serve  as  the  source  of  the  dioxide  used  in  the  process ;  otherwise  the  catalytic 
agent  loses  its  power.  (2)  An  excess  of  oxygen  over  that  required  in  equation  (2) 
must  be  present.  (3)  The  temperature  must  be  maintained  at  about  400°. 
(4)  In  order  that  all  the  sulfur  trioxide  formed  may  be  completely  absorbed, 
it  is  passed  into  sulf uric  acid  containing  from  2  to  3  per  cent  of  water.  As  the 
trioxide  is  absorbed  a  corresponding  amount  of  water  is  constantly  run  in,  so 
that  the  absorbing  solution  retains  a  constant  concentration. 

2.  Lead-chamber  process .  This  process  receives  its  name  from  the 
fact  that  the  reactions  are  carried  out  in  large  lead-lined  chambers. 
The  following  substances  enter  into  the  reaction :  (1)  sulfur  dioxide 
obtained  as  in  the  contact  process ;  (2)  a  mixture  of  nitric  oxide 
and  nitrogen  dioxide  obtained  by  heating  nitric  acid;  (3)  water  in 
the  form  of  steam ;  (4)  oxygen  (air). 

The  sulfides  of  some  of  the  metals,  such  as  zinc,  lead,  and  copper,  constitute 
the  ores  from  which  these  metals  are  largely  obtained.  In  the  extraction  of  the 
metals  from  these  ores  it  is  often  found  most  convenient  to  first  convert  the  sul- 
fides into  oxides  by  heating  them  in  the  air.  The  resulting  sulfur  dioxide  may 
be  utilized  in  the  manufacture  of  sulf  uric  acid.  It  often  happens,  therefore,  that 
the  manufacture  of  sulf  uric  acid  is  carried  on  in  connection  with  the  metallurgy 
of  these  metals. 

The  four  substances  mentioned  above,  namely,  sulfur  dioxide,  oxides 
of  nitrogen,  steam,  and  air,  are  introduced  in  the  proper  proportions 
into  large  lead-lined  chambers,  and  under  suitable  conditions  they  re- 
act to  form  sulfuric  acid.  The  reactions  involved  are  quite  complex, 
and  are  not  at  all  thoroughly  understood.  It  is  believed,  however,  that 
the  two  following  general  reactions  take  place :  (1)  The  substances 
introduced  into  the  chambers  first  react  to  form  a  derivative  of  sulfuric 


SULFUR;   SELENIUM;   TELLURIUM  219 

acid  known  as  nitrosyl-sulfuric  acid.    The  relation  of  these  two  com- 
pounds to  each  other  may  be  seen  from  their  structural  formulas  : 


HCT        O  NO-CT     ^O 

sulfuric  acid  nitrosyl-sulfuric  acid 

In  accordance  with  these  formulas  nitrosyl-sulfuric  acid  differs  from 
sulfuric  acid  in  composition  simply  by  containing  the  univalent  radical 
-  NO  (known  as  the  nitrosyl  group)  in  place  of  one  of  the  hydrogen 
atoms  of  sulfuric  acid.  Its  formation  may  be  represented  as  follows  : 


2  S02  +  NO  +  N02  +  H20  +  02  =  2  NO 

This  acid  can  be  obtained  in  the  form  of  white  crystals  known  as 
chamber  crystals. 

(2)  In  the  commercial  manufacture  of  sulfuric  acid,  however,  such 
a  separation  does  not  occur,  because  sufficient  water  is  always  present 
to  change  the  nitrosyl  acid,  as  fast  as  formed,  into  sulfuric  acid  : 


It  will  be  noted  that  in  this  second  reaction  the  same  quantities  of  the 
oxides  of  nitrogen  are  formed  as  are  required  for  the  first  reaction. 
Theoretically,  therefore,  a  small  amount  of  these  oxides  should  suffice 
to  prepare  an  unlimited  amount  of  sulfuric  acid  ;  practically,  some  of 
the  oxides  are  lost,  and  this  loss  must  be  replaced.  The  sulfuric  acid 
collects  upon  the  floor  of  the  chambers  in  the  form  of  an  aqueous 
solution  containing  from  62  to  70  per  cent  of  acid.  This  product  is 
called  chamber  acid  and  is  quite  impure,  but  for  many  purposes,  such 
as  the  manufacture  of  fertilizers,  it  needs  no  further  treatment.  It 
can  be  concentrated  by  evaporation  in  open  lead  pans  until  it  contains 
about  80  per  cent  acid.  Further  concentration  up  to  94  per  cent  may 
be  effected  in  platinum  vessels,  and  from  this  concentration  up  to  98 
per  cent  in  iron  vessels.  Prepared  in  this  way  the  ordinary  concen- 
trated commercial  sulfuric  acid  usually  contains  about  5  per  cent 
water  as  well  as  small  amounts  of  impurities,  especially  lead  dissolved 
from  the  chambers.  It  also  frequently  contains  traces  of  arsenic  origi- 
nally present  in  the  sulfides  used  as  the  source  of  the  sulfur  dioxide. 
These,  however,  do  not  interfere  with  its  use  for  many  purposes.  In 
order  to  obtain  the  chemically  pure  acid,  the  commercial  acid  is  dis- 
tilled from  glass  or  porcelain  vessels.  The  distillate  consists  of  approx- 
imately 98  per  cent  acid  and  2  per  cent  water. 


220 


GENERAL   CHEMISTRY 


The  sulfuric-acid  plant.  The  simpler  parts  of  a  plant  used  in  the  manufacture 
of  sulfuric  acid  are  illustrated  in  Fig.  86.  Sulfur,  or  some  sulfide,  as  FeS2,  is 
burned  in  the  furnace  A.  The  resulting  sulfur  dioxide,  together  with  the  neces- 
sary amount  of  air,  passes  into  the  structure  C,  known  as  the  Glover  tower.  In 
it  the  oxides  of  nitrogen  are  generated,  as  will  be  explained  below,  and  these, 
together  with  the  sulfur  dioxide  and  air,  pass  into  the  chambers  D,  D,  Water 
or  steam  is  also  introduced  into  these  chambers  at  suitable  points.  Here  the 
reactions  take  place  which  result  in  the  formation  of  the  sulfuric  acid.  The 
nitrogen  remaining  after  the  withdrawal  of  the  oxygen  from  the  air  which 
entered  the  chamber  escapes  through  the  structure  E,  known  as  the  Gay-Lussac 
tower.  In  order  to  prevent  the  escape  of  the  nitrogen  dioxide  regenerated  in  the 
reaction,  this  tower  is  filled  with  pieces  of  coke  over  which  trickles  concentrated 


FIG.  86 

sulfuric  acid  admitted  in  the  form  of  a  spray  (F)  at  the  top.  The  concentrated 
acid  absorbs  the  nitrogen  dioxide  but  not  the  nitric  oxide,  so  that  the  latter 
escapes  along  with  the  nitrogen.  The  acid  which  is  sprayed  into  the  top  of  the 
tower  collects  in  the  bottom  and  is  run  off  into  the  vessel  G,  from  which  it  is 
forced  into  the  tank  at  the  top  of  the  Glover  tower  C.  Here  it  is  mixed  with 
some  dilute  sulfuric  acid,  and  the  mixture  sprayed  into  the  top  of  the  tower, 
which  is  partly  filled  with  some  acid-resisting  rock.  As  the  acid  passes  down 
through  this  material  it  meets  with  the  hot  gases  entering  from  the  furnace, 
whereby  the  nitrogen  dioxide  is  liberated  from  the  acid,  passes  over  into  the 
chamber  D,  and  again  enters  into  the  reaction.  During  the  process  just  described 
the  dilute  acid  becomes  sufficiently  concentrated  to  serve  again  as  an  absorbent 
of  nitrogen  dioxide.  The  necessary  quantity  of  it  is  therefore  run  into  the  vessel 
H  from  the  bottom  of  the  tower,  and  then  forced  into  the  tank  at  the  top  of  E. 
In  order  to  replace  the  oxides  of  nitrogen  lost  in  the  process,  the  necessary  quan- 
tity is  added  by  the  action  of  sulfuric  acid  upon  sodium  nitrate  in  vessel  B.  The 
sulfuric  acid  formed  collects  in  the  bottom  of  the  chambers  and  is  drawn  off 
from  time  to  time. 

Historical.  The  chamber  process  is  a  very  old  one,  dating  back  to  the  year 
1746-  For  a  number  of  years  previous  to  this  date  sulfuric  acid  had  been  made 
by  burning  sulfur  mixed  with  potassium  nitrate  in  large  glass  globes.  Since 
these  globes  were  very  fragile,  and  since  lead  was  known  to  be  acted  upon  to  but 


SULFUR;   SELENIUM;   TELLURIUM  221 

a  very  slight  extent  by  the  dilute  acid,  Dr.  Roebuck,  an  Englishman,  suggested 
the  substitution  of  lead-lined  chambers  for  the  glass  globes.  The  first  plant  for 
the  manufacture  of  the  acid  by  this  method  was  constructed  in  Birmingham, 
England,  in  1746.  The  chambers  used  in  the  original  plant  were  very  small, 
having  a  capacity  of  about  200  cubic  feet,  while  some  of  those  in  use  at  the 
present  time  have  a  capacity  as  great  as  75,000  cubic  feet. 

Relative  advantages  of  the  contact  process  and  lead-chamber  process.  It  will  be 
noted  that  in  the  contact  process  it  is  just  as  easy  to  prepare  the  pure  concen- 
trated acid  as  the  dilute  acid.  In  the  chamber  process,  however,  the  dilute  acid 
is  obtained  first  and  can  be  prepared  at  a  very  low  cost.  The  concentration  and 
purification  of  the  dilute  acid  is,  however,  an  expensive  operation.  For  these 
reasons  the  contact  process  can  compete  with  the  lead-chamber  process  only  in 
the  manufacture  of  the  pure  concentrated  acid.  The  contact  process  is  also  adapted 
to  the  manufacture  of  fuming  sulfuric  acid  (p.  228). 

Properties.  Pure  anhydrous  sulfuric  acid,  more  properly  termed 
hydrogen  sulfate,  is  a  colorless,  oily  liquid.  Because  of  its  oily 
appearance,  together  with  the  fact  that  it  was  formerly  obtained  by 
distilling  the  sulfate  of  iron  known  as  green  vitriol,  it  is  often 
termed  oil  of  vitriol.  Its  density  is  1.838  at  15°.  When  heated  to 
338°  it  boils,  a  portion  of  the  compound  decomposing  into  sulfur 
trioxide  and  water  in  the  process.  The  sulfur  trioxide  escapes  to  a 
greater  extent  than  does  the  water,  until  the  residual  liquid  contains 
1.67  per  cent  of  water.  The  resulting  mixture  then  distills  with 
unchanged  concentration  (p.  127).  At  a  low  temperature,  hydrogen 
sulfate  forms  crystals  which  melt  at  10.5°. 

Chemical  conduct.  The  chemical  conduct  of  hydrogen  sulfate  may 
be  discussed  under  the  general  topics  enumerated  below. 

1.  Acid  properties.    In  concentrated  aqueous  solutions,   hydrogen 
sulfate  forms  the  ions  H+  and  HSO4~,  the  latter,  on  further  dilution  of 
the  solution,  breaking  down  into  the  ions  H+  and  SO4~~.    It  is  this 
aqueous  solution  containing  hydrogen  ions  which  is  properly  termed 
sulfuric  acid. 

2.  Dissociation  by  heat.    When  heated  above  100°,  dissociation  into 
water  and  sulfur  trioxide  becomes  marked  and  gradually  increases 
with  rising  temperature  until,  at  about  450°,  it  is  complete.    The  re- 
action is  reversible,  recombination  of  the  water  and  sulfur  trioxide 

taking  place  as  the  temperature  falls : 

/x     n  ,~r 

\~s  .  -  n  ^  y  P    ^  ^  " 

3.  Oxidizing  properties.  Hydrogen  sulfate  is  a  strong  oxidizing  agent, 
being  similar  in  this  respect  to  hydrogen  nitrate  (nitric  acid).    When 


222  GENERAL  CHEMISTRY 

heated  with  substances  which  readily  combine  with  oxygen,  such  as 
carbon,  sulfur,  and  many  of  the  metals,  it  gives  up  oxygen  to  these 
substances,  being  itself  reduced  to  sulf  urous  acid  : 


It  will  be  recalled  that  this  general  reaction  serves  as  one  of  the 
methods  for  preparing  sulfur  dioxide  (p.  212).  Because  of  its  strong 
oxidizing  action,  hydrogen  sulf  ate  is  used  commercially  in  the  prepa- 
ration of  a  number  of  compounds,  such  as  indigo,  and  has  an  advan- 
tage in  that  the  sulfur  dioxide  formed  in  the  process  can  be  changed 
back  into  hydrogen  sulfate  by  the  contact  method,  and  may  thus  be 
utilized  repeatedly. 

Hydrogen  sulfate  thus  resembles  its  anhydride  SO3  in  that  both  are 
strong  oxidizing  agents,  much  as  sulf  urous  acid  resembles  its  anhy- 
dride SO2  in  reducing  properties.  It  is  a  general  rule  that  an  acid  and 
its  anhydride  are  very  similar  in  oxidizing  and  reducing  properties, 
and  many  examples  will  be  found  in  subsequent  pages. 

4.  Action  upon  metals.  Hydrogen  sulfate  reacts  with  most  of  the 
metals,  especially  if  heated  in  contact  with  them.  The  action  is  simi- 
lar to  that  of  hydrogen  nitrate  in  that  the  metal  is  first  oxidized,  the 
resulting  oxide  being  then  changed  into  the  corresponding  salt.  Thus, 
the  reaction  between  copper  and  hydrogen  sulfate  takes  place  according 
to  the  following  equations  : 

Cu  +  H2SO4  =  CuO  +  H2SO8 


CuO  +  H2SO4  =  CuSO4  +  H2O 

Canceling  the  formulas  of  the  compounds  formed  in  one  reaction  and 
used  up  in  another,  —  namely,  H2SO3  and  CuO,  —  these  equations  may 
be  combined  into  the  following  : 

Cu  4-  2  H2SO4  =  CuSO4  4-  2  H2O  +  SO2 

The  conduct  of  a  dilute  solution  of  hydrogen  sulfate  (that  is,  dilute 
sulfuric  acid)  toward  metals  is  entirely  different  from  that  of  hydrogen 
sulfate  itself.  While  it  has  practically  no  action  upon  the  metals  occur- 
ring below  hydrogen  in  the  electromotive  series,  it  reacts  more  or  less 
readily  with  those  occurring  above  hydrogen,  evolving  hydrogen  and 
forming  the  corresponding  sulfates.  It  will  be  recalled  that  this  reac- 
tion serves  as  the  common  laboratory  method  for  preparing  hydrogen. 


SULFUR;   SELENIUM;  TELLURIUM  223 

But  few  of  the  metals  entirely  resist  the  action  of  hydrogen  sulfate. 
Even  platinum  and  gold,  which  are  very  resistant  to  the  action  of  acids, 
are  perceptibly  dissolved  when  heated  with  it.  While  iron  readily 
dissolves  in  the  dilute  acid,  it  is  only  slightly  attacked  by  the  more 
concentrated  acid,  so  that  iron  vessels  are  used  for  effecting  the  final 
concentration  in  the  preparation  of  sulfuric  acid. 

5.  Action  upon  salts.  The  action  of  hydrogen  sulfate  upon  sodium 
nitrate  (p.  200)  is  typical  of  its  action  upon  salts  of  all  acids  having 
a  lower  boiling  point  than  the  hydrogen  sulfate  itself.    This  action  has 
been  explained  in  detail  in  the  preceding  chapter  (p.  196). 

6.  Action  upon  water.    When  hydrogen  sulfate  is  mixed  with  water, 
a  marked  contraction  in  volume  takes  place.    Moreover,  the  process 
is  attended  by  the  evolution  of  a  large  amount  of  heat,  which  may 
even  be  sufficient  to  cause  the  liquid  to  boil.    Care  must  therefore 
be  taken,  in  mixing  the  sulfate  with  water,  to  keep  the  solution 
thoroughly  stirred  during  the  process,  and  to  pour  the  sulfate  into 
the  water,  not  the  reverse. 

When  hydrogen  sulfate  and  water  are  mixed  in  molecular  propor- 
tions—  namely,  98  parts  by  weight  of  the  sulfate  to  18  of  water  — 
and  the  resulting  solution  is  cooled,  the  two  unite  to  form  the  hydrate 
H2SO4  •  H2O.  This  is  a  white,  crystalline  substance  melting  at  8.5°. 

7.  Action  upon  organic  compounds.   Not  only  does  hydrogen  sulfate 
absorb  water  itself,  but  it  decomposes  many  compounds  containing 
the  elements  hydrogen  and  oxygen.   When  acted  upon  by  hydrogen 
sulfate,  the  hydrogen  and  oxygen  present  in  such  compounds  unite  to 
form  water,  which  is  then  absorbed  by  the  sulfate.    For  example,  most 
organic  substances,  such  as  sugar,  wood,  cotton  and  woolen  fiber,  and 
even  flesh  (all  of  which  contain  hydrogen  and  oxygen  in  addition  to 
carbon),  are  charred  by  the  action  of  the  sulfate,  the  charring  being  due 
to  the  withdrawal  of  the  hydrogen  and  oxygen  from  the  compound. 

Structural  formula  of  hydrogen  sulfate.  The  chemical  conduct  of 
hydrogen  sulfate  is  best  explained  upon  the  assumption  that  the  mole- 
cule contains  two  hydroxyl  groups,  and  that  these  groups,  as  well  as  the 
remaining  oxygen  atoms,  are  directly  bound  to  the  sulfur,  as  indicated 
in  the  following  formula,  in  which  the  sulfur  is  hexavalent : 

HO.  <^0 

Her  ^o 
This  is  often  abbreviated  to  the  form  SO2(OH)2. 


224 


GENEKAL  CHEMISTRY 


Salts  of  sulfuric  acid ;  sulfates.  The  sulfates  constitute  a  very  im- 
portant class  of  compounds,  and  many  of  them  have  extensive  com- 
mercial uses.  The  normal  salts  are  all  solids  and,  with  the  exception 
of  those  of  barium,  strontium,  and  lead,  are  soluble  in  water.  Two 
others,  namely,  calcium  sulfate  and  silver  sulfate,  are  only  slightly 
soluble.  The  hydrates  of  many  of  the  sulfates  are  more  frequently 
used  than  the  anhydrous  salts.  Some  of  the  more  important  of  these 
compounds  are  included  in  the  following  table : 


FORMULA 


CHEMICAL  NAME 


COMMON  NAME 


Na2SO4  •  10  H2O 
MgSO4  •  7  H2O 
CaS04  •  2  H20 
CuS04  •  5  H20 
FeSO4  •  7  H2O 
ZnSO4 
BaSO4 


7H2O 


Hydrate  of  sodium  sulfate 
Hydrate  of  magnesium  sulfate 
Hydrate  of  calcium  sulfate 
Hydrate  of  copper  sulfate 
Hydrate  of  iron  sulfate 
Hydrate  of  zinc  sulfate 
Barium  sulfate 


Glauber's  salt 

Epsom  salts 

Gypsum 

Blue  vitriol 

Green  vitriol  or  copperas 

White  vitriol 

Barite  or  heavy  spar 


The  sulfates  can  be  prepared  by  the  action  of  sulfuric  acid  upon 
the  metals,  their  oxides  and  hydroxides,  and  many  of  their  salts,  as 
illustrated  by  the  following  equations  : 


Zn(OH)2  +  H2SO4  =  ZnSO4  +  2  H2O 

2  HNO 
H2S 


2  NaN03  +  H2SO4  =  Na2SO4 
FeS  +  H2SO4  =  FeSO4 


The  insoluble  sulfates  may  be  prepared  by  the  general  method  used 
in  preparing  insoluble  compounds.  Thus,  barium  chloride  (BaCla)  and 
sulfuric  acid  in  aqueous  solutions  react  as  indicated  in  the  equation  : 


Ba 


,  SO4~  -  -  BaSO4 


The  resulting  barium  sulfate  separates  in  the  form  of  a  white  solid. 
Since  only  soluble  sulfates  and  sulfuric  acid  yield  the  ion  SO4~~  in 
solution,  this  reaction  serves  as  a  delicate  test  for  these  compounds. 

The  action  of  water  upon  normal  salts.  If  a  dilute  solution  of  sodium 
sulfate  in  water  is  tested  with  litmus,  it  is  found  to  be  neutral.  One 
might  naturally  expect  solutions  of  all  normal  salts  to  act  in  the  same 
way.  In  reality  they  do  not  do  so.  For  example,  a  solution  of  sodium 
sulfide  (Na2S)  or  sodium  sulfite  (Na2SO3)  is  distinctly  basic  in  reac- 
tion, while  a  solution  of  copper  sulfate  (CuSO4)  or  aluminium  sulfate 


SULFUR;   SELENIUM;  TELLUEIUM  225 

(A12(SO4)3)  is  acid.  This  difference  in  the  reaction  of  the  aqueous 
solutions  of  different  classes  of  salts  is  in  accord  with  the  ionization 
theory,  as  will  be  clear  from  the  following  discussion. 

It  will  be  recalled  that  water  itself  is  slightly  ionized  as  follows  : 

H20  +±.  H+,  OH- 

If  some  substance  is  added  to  the  water  which  will  continuously  with- 
draw either  the  hydrogen  or  the  hydroxyl  ions,  then  more  water  will 
be  ionized,  in  order  that  the  equilibrium  expressed  in  the  equation 
may  be  maintained.  If  this  action  continues,  the  concentration  of  the 
ion  not  withdrawn  may  be  increased  to  such  an  extent  that  its  effect 
will  become  marked.  The  resulting  solution  will  therefore  react  acid 
or  basic  according  to  whether  it  is  the  hydroxyl  or  the  hydrogen  ion 
which  has  been  withdrawn. 

In  discussing  this  general  subject  it  is  convenient  to  divide  the  salts 
into  four  general  classes,  as  designated  below. 

1.  Salts  derived  from  strong  bases  and  weak  acids.  All  salts  belonging 
to  this  class  give  basic  solutions.  Sodium  sulflde  will  serve  as  an 
example,  being  derived  from  sodium  hydroxide  and  hydrosulfuric  acid. 
When  this  salt  is  added  to  water,  the  reactions  expressed  in  the 
following  equations  take  place  : 

=^Na+,  Na+,  S-- 


H2O 


But  each  positive  ion  in  the  solution  must  be  in  equilibrium  with 
each  negative  ion,  so  that  there  will  be  two  other  equilibria: 


Na+,  OH-  +=±  NaOH 

In  the  first  of  these  the  product  formed,  H2S,  is  a  very  weak  acid  and 
is  consequently  little  ionized,  so  that  equilibrium  is  not  reached  until 
almost  all  of  the  hydrogen  ions  have  entered  into  combination.  In  the 
second  equilibrium  the  product  formed,  NaOH,  is  a  strong  base  and 
is  largely  ionized.  As  a  result  of  these  differences  more  hydrogen  ions 
are  withdrawn  in  the  first  equilibrium  than  hydroxyl  ions  in  the  second, 
and  the  solution  acquires  strong  basic  properties  from  the  excess  of 
hydroxyl  ions. 

2.  Salts  derived  from  strong  acids  and  weak  bases.   Salts  belonging  to 
this  class  yield  acid  solutions.  Copper  sulfate  will  serve  as  an  example. 


226  GENERAL  CHEMISTRY 

When  this  salt  is  dissolved  in  water,  the  equilibria  represented  in  the 
following  equations  result : 

CuSO4:*=^Cu++,  SO4~- 
H2O  +=±  H+,  OH- 

But  in  addition  to  these  there  must  be  the  further  equilibria  ex- 
pressed in  the  following  equations : 

Cu++,  OH-,  OH-  +=±  Cu(OH)2 
H+,  H+,  S04-^H2S04 

In  the  first  of  these  the  product  formed,  Cu(OH)2,  is  a  very  weak 
base  and  is  therefore  little  ionized,  so  that  at  equilibrium  most  of  the 
hydroxyl  ions  are  withdrawn  from  the  solution.  With  sulfuric  acid 
formed  in  the  second  equilibrium  the  case  is  quite  different,  since 
it  is  a  strong  acid  and  is  largely  ionized.  As  a  result  of  these 
differences  more  hydroxyl  ions  are  withdrawn  than  hydrogen  ions, 
and  consequently  the  solution  acquires  acid  properties. 

3.  Salts  derived  from  strong  bases  and  strong  acids.   All  such  salts 
yield  neutral  solutions.    It  is  evident  from  the  above  discussions  that 
when  salts  of  this  class  are  dissolved  in  water,  neither  the  hydrogen 
nor  the  hydroxyl  ions  are  withdrawn  in  appreciable  amounts ;  hence 
their  solutions  are  neutral. 

4.  Salts  derived  from  weak  bases  and  weak  acids.    Such  salts  are  more 
or  less  completely  decomposed  in  the  presence  of  water. 

Hydrolysis.  If  we  wish  simply  to  represent  the  compounds  reacting 
and  the  products  formed,  the  reactions  taking  place  when  such  salts 
as  sodium  sulfide  and  copper  sulfate  are  dissolved  in  water  may  be 
expressed  as  follows: 

Na2S  +  2  HOH  +=±  2  NaOH  +  H2S 
CuSO4+  2  HOH  +=£  Cu(OH)2+  H2SO4 

It  will  be  noted  that  these  reactions  belong  to  the  general  type 
designated  by  the  term  double  decomposition  (p.  100),  the  metal  of 
the  salt  changing  places  with  the  hydrogen  of  the  water.  A  double 
decomposition  in  which  water  is  one  of  the  reacting  compounds  is 
termed  hydrolysis.  The  compound  reacting  with  the  water  is  said  to 
undergo  hydrolysis.  In  each  of  the  cases  just  described  the  hydrolysis 
is  only  partial,  since  the  equilibrium  expressed  in  the  equation  is  soon 
reached.  We  shall  meet  with  other  examples,  however,  in  which  the 
hydrolysis  is  practically  complete. 


SULFUR;  SELENIUM;   TELLURIUM  227 

The  action  of  water  upon  acid  salts.  The  acid  salts  of  strong  acids 
readily  ionize  in  aqueous  solutions  yielding  hydrogen  ions.  Thus, 
sodium  hydrogen,  su  If  ate  ionizes  as  follows: 

NaHS04  :<=±  Na+,  HS<V  (1) 

^=±H+,  S04-  (2) 


In  dilute  solutions  both  of  the  reactions  expressed  in  (1)  and  (2)  are 
practically  complete.  Such  solutions,  therefore,  have  an  acid  reaction 
due  to  the  hydrogen  ions.  The  case  is  different,  however,  with  acid 
salts  of  weak  acids,  since  the  hydrogen  present  in  such  salts  is  not 
ionized  to  any  appreciable  extent  when  the  salts  are  dissolved  in 
water.  Sodium  hydrogen  sulfide,  for  example,  ionizes  as  follows  : 

NaHS  +=±:  Na+,  HS~  (1) 

HS-:<=±H+,  S--  (2) 

The  reaction  expressed  in  (2),  however,  takes  place  only  to  a  very 
slight  extent.  A  solution  like  this  will  therefore  contain  such  a  com- 
paratively small  number  of  hydrogen  ions  that  it  will  have  either  a 
neutral  or,  at  most,  only  a  slightly  acid  reaction.  Indeed,  in  the  case 
of  acid  salts  derived  from  a  strong  base  and  a  very  weak  acid,  hydrol- 
ysis may  take  place  to  such  an  extent  that  their  solutions  react  basic. 
For  example,  sodium  hydrogen  carbonate  (NaHCO3),  an  acid  salt  of 
carbonic  acid  (H2CO3),  dissolves  in  water,  forming  the  ions  Na+  and 
HCO3~.  Since  carbonic  acid  is  a  very  weak  acid  scarcely  ionizing  at 
all  in  solution,  it  will  be  formed  by  the  union  of  ions  HCO3~  with  the 
hydrogen  ions  of  the  water  : 

H+  +  HC03-  +=*  H2C03 

This  withdrawal  of  the  hydrogen  ions  from  the  water  results  in  the 
accumulation  of  the  accompanying  hydroxyl  ions  to  such  an  extent 
that  the  solution  becomes  basic  in  reaction. 

Other  oxygen  acids  of  sulfur.  In  addition  to  sulfurous  and  sulfuric 
acids  a  number  of  other  oxygen  acids  of  sulfur  are  known,  either  in 
the  free  state  or  in  the  form  of  their  salts.  The  formulas  and  names 
of  these  are  as  follows  : 

Thiosulfuric  acid        .     .     .  H2S2O3  Dithionic  acid  .....  H2S2O6 

Hyposulfurous  acid    .     .     .  H2S2O4  Trithionic  acid       ....  H2S3O6 

Pyrosulfuric  acid        '.     .     .  H2S2O7  Tetrathionic  acid  ....  H2S4O6 

Persulfuric  acid      ....  H2S2O8  PentatMonic  acid        .     .     .  H2S5O6 


228  GENERAL  CHEMISTRY 

The  thionic  acids  derive  their  names  from  a  Greek  word  meaning 
"sulfur."  With  the  exception  of  pyrosulfuric  acid,  the  acids  are  so 
unstable  that  they  have  been  obtained  only  in  dilute  aqueous  solution 
or  in  the  form  of  salts.  A  brief  discussion  of  pyrosulfuric  and  per- 
sulfuric  acid  follows.  The  important  salts  of  the  other  acids  will  be 
discussed  in  their  appropriate  places  in  connection  with  the  metals. 

Pyrosulfuric  acid  (disulfuric  acid)  (H2S207).  When  sulfuric  acid  and 
sulfur  trioxide  are  brought  together  in  molecular  proportions,  there  is 
formed  a  solid  crystalline  compound  known  as  pyrosulfuric  acid  : 


This  acid  finds  some  use  as  a  powerful  oxidizing  agent.  Its  salts  can 
be  prepared  by  heating  the  corresponding  acid  sulfates.  The  hydro- 
gen present  in  the  salt  is  evolved  in  the  form  of  water,  leaving 
the  pyrosulfate: 

2  NaHS04  =  Na2S207  +  H2O 

The  fuming  sulfuric  acid  of  commerce  consists  of  sulfuric  acid  con- 
taining various  percentages  of  pyrosulfuric  acid. 

Persulfuric  acid  (H2S208).  It  will  be  recalled  that  in  concentrated 
solutions  of  sulfuric  acid  there  exists  the  equilibrium  expressed  in 
the  following  equation: 

=±:H+,  HS0- 


4 


When  such  solutions  are  electrolyzed,  hydrogen  is  evolved  at  the 
cathode.  At  the  anode  the  ions  HSO4~  are  discharged,  and  the  result- 
ing radicals  combine  with  each  other  to  form  the  compound  (HSO4)2 
or  H2S2Og  which  remains  dissolved  in  the  liquid  about  the  anode. 
This  compound  is  known  as  persulfuric  acid.  While  the  acid  is  un- 
stable and  exists  only  in  dilute  solution,  its  salts  can  be  obtained  in 
pure  condition.  They  are  prepared  by  the  electrolysis  of  concentrated 
solutions  of  the  corresponding  acid  sulfates.  For  example,  a  concen- 
trated solution  of  potassium  hydrogen  sulfate,  upon  electrolysis,  yields 
persulfuric  acid  at  the  anode,  just  as  does  the  electrolysis  of  sulfuric 
acid  itself.  The  acid,  however,  as  fast  as  formed,  reacts  with  the 
potassium  hydrogen  sulfate  present  in  the  solution  to  form  potassium 
persulfate  : 

2  KHS04  +  H2S208  =  K2S208 


The  potassium  salt,  being  sparingly  soluble,  crystallizes  as  a  white  solid. 


SULFUR;   SELENIUM;  TELLURIUM  229 

Persulfuric  acid,  as  well  as  its  salts,  is  a  strong  oxidizing  agent 
In  the  presence  of  a  reducing  agent,  aqueous  solutions  of  the  acid 
and  its  salts  decompose  as  follows,  the  oxygen  combining  with  the 

reducing  agent: 

2  H2S208  +  2  H20  =  4  H2S04  +  2  O 
2  K2S208  +  2  H20  =  4  KHS04  +  2  O 

Some  of  the  persulfates  are  coming  into  use  as  commercial  oxidizing 
agents.  For  example,  ammonium  persulfate  is  being  used  in  certain 
photographic  processes,  as  well  as  for  a  general  oxidizing  agent. 

Compounds  containing  sulfur  and  chlorine.  The  following  compounds 
containing  sulfur  and  chlorine  are  known  : 

1.  Sulfur  monochloride  («S2C/2).  This  compound  was  formerly  thought 
to  have  the  formula  SCI,  hence  the  name  monochloride.    It  is  formed 
when  chlorine  is  passed  over  heated  sulfur,  and  is  a  yellow,  oily  liquid 
boiling  at  138°  and  having  an  extremely  disagreeable  odor.  It  is  a  good 
solvent  for  sulfur,  and  this  solution  is  used  in  vulcanizing  rubber. 

2.  Sulfur  tetrachloride  (SC/J.    At  low  temperatures,  sulfur  mono- 
chloride  combines  with  additional  chlorine  to  form  a  reddish-brown 
liquid  which  has  the  composition  SC14.    This  compound  is  stable  only 
at  temperatures  below  —  20°.    As  the  temperature  rises,  it  gradually 
decomposes  into  sulfur  monochloride  and  free  chlorine. 

3.  Thionyl  chloride  (SOC/J.    This  compound  is  obtained,  along  with 
phosphorus  oxy  chloride  (POC13),  by  the  action  of  sulfur  dioxide  upon 
phosphorus  pentachloride  (PC15)  : 

SO2  +  PC15  =  SOC12  +  POC18 

Thionyl  chloride  is  a  colorless  liquid  boiling  at  78°.  With  water  it 
decomposes  as  indicated  in  the  following  equation  : 


4.  Sulfuryl  chloride  (S02C/3).  Under  proper  conditions  sulfur  dioxide 
and  chlorine  combine  directly  to  form  sulfuryl  chloride,  a  compound 
having  the  formula  SO0C12.  The  reaction  may  be  greatly  hastened  by 
using  camphor  as  a  catalytic  agent.  Sulfuryl  chloride  is  a  colorless 
liquid  and  has  a  pungent  odor.  It  has  a  density  of  1.67  and  boils 
at  69°.  In  the  presence  of  a  relatively  small  quantity  of  water  one 
of  the  chlorine  atoms  is  displaced  by  a  hydroxyl  group,  forming  the 
compound  SO2C1(OH),  known  as  chlorosulfonic  acid  : 

SQ2C12  +  HOH  =  SO2C1(OH)  +  HC1 


230  GENERAL  CHEMISTRY 

With  an  excess  of  water  both  atoms  of   chlorine  are  displaced  by 
hydroxyl  groups,  forming  sulfuric  acid: 

S02C12  +  2  HOH  =  S02(OH)2  +  2  HC1 

The  structural  relations  of  sulfuric  acid,  chlorosulfonic  acid,  and  sul- 
furyl  chloride  to  each  other  are  shown  in  the  following  formulas : 


U0^»vi  w  C1< 

The  valence  of  sulfur.  The  valence  of  sulfur  varies  in  its  different 
compounds.  It  is  reasonably  certain  that  it  is  divalent  in  hydrogen 
sulfide  H— S— H  and  hexavalent  in  hydrogen  sulfate.  In  sulfurous 
acid  it  is  uncertain  whether  it  is  tetravalent  or  hexavalent  (see 
p.  215).  If  the  oxygen  atoms  in  sulfur  dioxide  and  sulfur  trioxide  are 
all  directly  bound  to  sulfur,  then  the  element  in  these  compounds  is 
tetravalent  and  hexavalent  respectively.  Nothing  is  known  of  the 
valence  of  sulfur  in  its  more  complex  acids  and  salts,  since  their 
structure  has  not  been  determined. 

SELENIUM  AND  TELLURIUM 

The  elements  selenium  and  tellurium  may  appropriately  be  consid- 
ered in  connection  with  sulfur,  because  of  their  intimate  relation  to 
this  element.  Selenium  resembles  sulfur  in  many  of  its  properties,  as 
well  as  in  its  chemical  conduct.  While  tellurium  does  not  show  this 
marked  similarity,  nevertheless  its  compounds  are  closely  related  in 
composition  to  those  of  sulfur. 

Tellurium  was  discovered  by  Reichenstein  and  Klaproth  in  the  lat- 
ter part  of  the  eighteenth  century.  The  latter  investigator  gave  to 
the  element  the  name  it  now  bears,  the  word  meaning  "the  earth." 
A  few  years  later  (1817)  Berzelius  isolated  a  new  element,  which  he 
named  selenium,  from  a  Greek  word  meaning  "  the  moon."  Both  of 
these  elements  occur  in  nature  in  the  free  as  well  as  in  the  combined 
condition,  but  only  in  comparatively  small  quantities,  being  far  less 
abundant  than  sulfur. 

Selenium.  This  element  is  frequently  found  in  small  quantities  in 
natural  sulfur.  Combined  with  metals  it  also  occurs  along  with  some 
of  the  sulfides,  especially  pyrites  (FeS2).  When  sulfur  or  a  sulfide 
containing  selenium  is  used  in  the  manufacture  of  sulfuric  acid,  some 


SULFUR;   SELENIUM;  TELLUEIUM  231 

free  selenium  is  always  found  in  the  flue  dust,  as  well  as  in  the 
slime  collecting  in  the  bottom  of  the  lead  chambers,  and  it  was  from 
this  material  that  Berzelius  first  isolated  the  element. 

Properties.  Selenium  resembles  sulfur  in  that  it  occurs  in  a  number 
of  different  forms,  although  the  relation  of  these  to  each  other  has  not 
been  so  well  determined  as  in  the  case  of  sulfur.  The  most  important 
of  these  forms  are  the  following :  (1)  a  red,  amorphous  solid,  slightly 
soluble  in  carbon  disulfide ;  (2)  a  red,  crystalline  solid,  likewise 
slightly  soluble  in  carbon  disulfide;  (3)  a  gray,  metallic,  crystalline 
solid,  insoluble  in  carbon  disulfide.  In  this  last  form  selenium  is  a 
conductor  of  electricity,  and  it  is  an  interesting  fact  that  its  conduc- 
tivity increases  with  the  intensity  of  the  light  to  which  the  selenium  is 
subjected.  When  heated  in  the  presence  of  air,  it  forms  selenium 
dioxide  (SeO2). 

Several  tons  of  selenium  are  recovered  each  year  in  the  United 
States  as  a  by-product  in  the  refining  of  copper,  and  more  could  be 
produced  if  there  were  a  demand  for  it.  It  is  used  almost  entirely 
in  the  manufacture  of  glass  and  enamels. 

Compounds  of  selenium.  In  general  it  may  be  stated  that  its  com- 
pounds resemble  those  of  sulfur,  in  composition  as  well  as  in  general 
properties.  The  most  important  of  these  compounds  are  as  follows : 
(1)  hydrogen  selenide  (H2Se),  a  gas  of  unbearable  odor;  (2)  sele- 
nium dioxide  (SeO2),  a  white,  crystalline  solid  formed  by  burning 
selenium  in  air ;  (3)  hydrogen  selenite  (H2SeO3),  an  unstable,  white, 
crystalline  solid  formed  by  the  combination  of  selenium  dioxide  and 
water;  (4)  hydrogen  selenate  (H2SeO4),  a  solid  melting  at  58°,  its 
aqueous  solutions  (selenic  acid)  resembling  sulfuric  acid  in  properties 
and  chemical  conduct ;  (5)  salts  derived  from  the  acids  HjSe,  H2SeO3, 
and  H2SeO4,  in  a  general  way  resembling  those  of  the  corresponding 
sulfur  compounds. 

Tellurium.  While  small  quantities  of  tellurium  occur  in  nature  in 
the  free  state,  it  is  much  more  frequently  found  combined  with  metals, 
especially  gold,  silver,  lead,  and  bismuth,  forming  compounds  known 
as  tellurides.  In  separating  these  metals  from  their  ores,  considerable 
quantities  of  tellurium  are  often  obtained,  although  no  commercial 
use  has  as  yet  been  found  for  it.  The  pure  element  resembles  the 
metals  in  appearance.  It  has  a  silvery  luster  and,  like  the  metals,  it 
conducts  electricity.  It  has  a  density  of  6.2  and  melts  at  450°.  When 
heated  in  the  air,  it  burns,  forming  tellurium  dioxide  ( 


232  GENERAL  CHEMISTEY 

Compounds  of  tellurium.  Some  of  the  more  important  compounds  of 
tellurium  are  the  following :  (1)  hydrogen  telluride  (H2Te),  a  gas  of 
disagreeable  odor,  which  resembles  hydrogen  sulfide  in  its  properties ; 
(2)  tellurium  dioxide  (TeO2),  a  white  solid  formed  by  the  combus- 
tion of  tellurium  in  oxygen  or  air;  (3)  tellurium  trioxide  (TeO3), 
an  orange-yellow  solid  formed  by  the  careful  heating  of  hydrogen  tel- 
lurate ;  (4)  hydrogen  tellurite  (H2TeO3),  a  white  powder  formed  by 
the  oxidation  of  tellurium  with  nitric  acid;  (5)  hydrogen  tellurate 
(H2TeO4),  a  solid  formed  by  the  action  of  strong  oxidizing  agents  upon 
tellurium,  its  aqueous  solution  (telluric  acid)  having  very  weak  acid 
properties.  Salts  of  the  different  acids  of  tellurium  are  also  known. 


CHAPTER  XVIII 

CLASSIFICATION  OF  THE  ELEMENTS 

Introduction.  Four  of  the  elements  —  namely,  oxygen,  hydrogen, 
nitrogen,  and  sulfur  —  have  now  been  studied  in  detail,  while  two 
others,  selenium  and  tellurium,  have  been  considered  more  briefly.  Of 
these,  oxygen,  hydrogen,  and  nitrogen  resemble  each  other  in  general 
properties,  while  in  chemical  conduct  they  are  quite  diverse.  Oxygen 
and  sulfur,  while  far  apart  in  physical  properties,  have  much  in  common 
in  the  types  of  compounds  which  they  form.  Sulfur,  selenium,  and 
tellurium  are  closely  related  both  in  physical  properties  and  in  chem- 
ical conduct,  and  in  most  respects  show  a  regular  gradation  in  prop- 
erties as  we  pass  from  sulfur  to  selenium  and  thence  to  tellurium. 

There  are  at  present  about  eighty  substances  classified  as  elements. 
It  is  evident  that  if  these  can  be  arranged  in  groups  in  which  a  some- 
what regular  gradation  in  properties  occurs,  the  study  of  the  elements 
will  be  simplified.  Moreover,  it  would  seem  to  be  rather  probable  that 
among  so  many  independent  substances  some  natural  relationships 
should  exist,  which,  if  discovered,  would  serve  as  the  best  basis  for 
classification.  Accordingly,  from  a  comparatively  early  time  chemists 
have  attempted  classifications  of  various  kinds,  at  first  basing  them 
on  more  obvious  chemical  relationships  and  gradually  seeking  a  more 
fundamental  and  natural  basis. 

Classification  based  on  chemical  conduct.  As  soon  as  a  reasonably 
clear  distinction  had  been  made  between  acids  and  bases  the  ele- 
ments were  divided  into  two  great  classes,  according  to  whether  their 
oxides  formed  acids  or  bases  on  hydration,  and  called  respectively  acid- 
forming  and  base-forming  elements.  Even  such  a  broad  classification 
was  not  free  from  difficulties,  since  many  elements  form  several  oxides, 
and  it  frequently  happens,  as  with  manganese  and  chromium,  that  some 
oxides  form  acids,  while  others  of  the  same  element  form  bases. 

Metals,  nonmetals,  and  metalloids.  Those  elements  which  usually 
form  bases  have  in  general  the  properties  which  we  usually  associate 
with  the  term  metal.  They  are  rather  heavy,  have  a  bright,  shining 
surface,  or  a  luster,  are  ductile  and  malleable,  and  are  good  conductors 


234  GENERAL  CHEMISTRY 

of  heat  and  electricity.  The  acid-forming  elements  have  the  opposite 
properties,  being  of  small  density,  of  little  or  no  luster,  brittle,  and  of 
small  conducting  capacity.  There  are,  however,  numerous  cases  in 
which  these  distinctive  marks  fail.  Thus,  sodium  and  potassium  are 
very  light,  though  undoubtedly  metals ;  silicon  is  rather  heavy  and 
has  a  high  luster,  though  a  iionmetal.  Gradually  the  terms  metal  and 
nonmetal  came  to  be  used  as  practically  synonymous  with  base-forming 
element  and  acid-forming  element.  An  intermediate  group  of  elements, 
which  possessed  in  some  degree  the  physical  properties  of  metals  and 
the  chemical  conduct  of  nonmetals,  came  to  be  known  as  metalloids. 
Examples  of  this  class  are  arsenic,  antimony,  and  tellurium. 

Classifications  based  on  atomic  weights.  Many  attempts  have  been 
made  to  base  a  classification  of  the  elements  upon  some  relation  be- 
tween their  atomic  weights.  These  are  the  most  characteristic  and 
fundamental  property  of  the  elements,  and  it  is  here,  if  anywhere,  that 
we  should  expect  to  find  a  basis  for  a  natural  grouping. 

Prout's  hypothesis.  As  early  as  1815,  upon  the  basis  of  very  imper- 
fect experiment,  Prout,  an  English  physician,  advanced  the  idea  that 
the  atomic  weights  of  all  the  elements  are  multiples  of  hydrogen 
taken  as  unity,  and  that  the  atoms  of  the  various  elements  consist  of 
varying  numbers  of  hydrogen  atoms.  Although  this  conception  has 
never  been  in  accord  with  the  best  experimental  work  of  the  time,  it 
has  possessed  a  fascination  for  many  minds,  and  has  stimulated  much 
of  the  best  earlier  work  upon  the  determination  of  atomic  weights. 

The  triads  of  Db'bereiner.  In  1829  Dobereiner  showed  that  many  of 
the  elements  may  be  arranged  in  groups  of  three  closely  similar  ones, 
in  which  the  atomic  weight  and  the  general  properties  of  the  one  is 
almost  an  exact  arithmetical  mean  between  the  other  two.  Thus  we  have 

Ca     .     .       40.07  Mg   .     .       24.32  Cl     .     .       35.46 

Sr      .     .       87.63  (88.72)      Zn    .     .       65.37  (68.36)      Br     . .   .       79.92  (81.19) 

Ba    .     .     137.37  Cd    .     .     112.40  I  ...     126.92 

A  number  in  parentheses  indicates  the  real  mean  between  the  weights. 

Many  other  relations  of  the  same  general  kind  were  discovered,  so 
that  chemists  came  to  feel  that  in  some  way  the  magnitude  of  the 
atomic  weight  of  an  element  really  determines  its  properties. 

The  periodic  grouping.  In  1869  two  distinguished  chemists,  the  Rus- 
sian MendeleefT  and  the  German  Lothar  Meyer,  quite  independently 
of  each  other,  discovered  a  relation  between  the  atomic  weights  of  the 


CLASSIFICATION  OF  THE  ELEMENTS  235 

elements,  which  has  come  to  be  known  as  the  periodic  law.  An  exact 
statement  of  the  law  will  be  better  understood  after  some  details  have 
been  explained. 

Plan  of  arrangement.  If  the  elements  are  arranged  in  the  order  of 
their  atomic  weights,  omitting  hydrogen,  the  first  eight  are : 

He  (3.99)    Li  (6.94)     Gl(9.1)    B(ll.O)     C  (12.0)    N  (14.01)    O(16.0)    F(19.0) 

These  elements  all  differ  markedly  from  each  other,  but  the  ninth 
element,  neon,  is  very  similar  to  helium.    It  is  placed  just  below  the 
latter,  and  a  new  row  follows : 
Ne(20.2)  Na(23.0)  Mg  (24.32)  Al(27.1)  Si  (28.3)  P  (31.04)  8(32.07)  Cl  (35.46) 

The  next  element,  argon,  again  resembles  helium  and  neon,  and  begins 
a  third  row : 

A  (39.88)  K(39.1)  Ca  (40.07)  Sc(44.1)  Ti(48.1)  V(51.0)  Cr(52.0)  Mn  (54.93) 

An  inspection  of  the  elements  in  the  eight  vertical  columns  brings 
to  light  a  remarkable  fact.  Not  only  are  helium,  neon,  and  argon  very 
similar,  but  a  more  or  less  pronounced  resemblance  is  found  between 
the  several  elements  in  each  of  the  columns.  Thus,  lithium,  sodium, 
and  potassium  are  very  similar,  as  are  glucinum,  magnesium,  and  cal- 
cium, as  well  as  carbon,  silicon,  and  titanium. 

The  three  elements  following  manganese  —  namely,  iron,  cobalt,  and 
nickel  —  resemble  each  other  very  closely,  and  if  arranged  as  the  first 
three  elements  in  a  fourth  row,  they  would  interrupt  the  regularities 
so  far  exhibited.  They  are  set  aside  in  a  ninth  column,  and  the  other 
elements  are  arranged  as  before.  The  complete  table  is  shown  on 
page  237.  A  vertical  row  is  called  a  group,  a  horizontal  row  a  series. 
It  will  be  noticed  that  two  other  sets  of  three  elements  are  placed 
in  the  ninth  group,  and  that  some  blank  spaces  appear  in  the  table, 
representing  elements  as  yet  undiscovered. 

The  relation  of  properties  to  atomic  weights.  There  is  evidently 
an  intimate  relationship  between  the  properties  of  an  element  and  its 
atomic  weight.  Helium,  at  the  beginning  of  the  first  row,  has  no 
chemical  activity  at  all.  Lithium  is  a  metallic  element  of  very  strong 
base-forming  character  and  a  valence  of  1.  Glucinum  is  also  a  metal, 
but  with  less  strongly  marked  base-forming  character  and  a  valence 
of  2.  Boron  is  a  metalloid,  is  much  more  acid-forming  than  basic  in 
character,  and  has  a  valence  of  3.  In  carbon,  with  a  valence  of  4,  all 
basic  properties  have  disappeared,  and  it  is  an  acid-forming  element 


236  GENERAL  CHEMISTRY 

exclusively.  The  succeeding  elements,  oxygen,  nitrogen,  and  fluorine, 
are  increasingly  acid-forming,  while  the  valence  increases  to  7.  A 
somewhat  similar  change  may  be  noticed  in  many  of  the  physical 
properties  of  the  elements,  such  as  their  conductivity  toward  heat 
and  electricity,  their  densities,  and  their  melting  points.  These  prop- 
erties usually  vary  with  some  regularity  as  we  pass  from  one  end 
of  a  series  to  the  other.  The  properties  of  these  elements,  there- 
fore, vary  more  or  less  regularly  with  their  atomic  weights,  or,  in 
mathematical  language,  the  properties  are  continuous  functions  of 
the  atomic  weights. 

The  periodic  law.  If  helium  were  the  element  of  smallest  atomic 
weight  and  fluorine  that  of  the  greatest,  so  that  in  passing  from  helium 
to  fluorine  we  had  passed  all  the  elements  in  review,  we  could  make 
the  general  statement  that  the  properties  of  the  elements  vary  directly 
with  their  atomic  weights.  But  fluorine  has  a  small  atomic  weight, 
and  neon,  which  follows  it,  repeats  the  properties  of  helium,  starting 
a  new  series,  or  period.  If  we  omit  the  elements  in  the  ninth  column, 
we  see  that  each  ninth  element  starts  a  new  period,  and  we  reach  a 
statement  of  the  periodic  law,  which  is:  The  properties  of  an  element 
are  periodic  functions  of  the  atomic  weight. 

The  atmospheric  elements.  Until  a  few  years  ago  the  elements  in 
the  column  marked  O  were  unknown.  When  they  were  discovered 
they  were  seen  to  constitute  a  column  by  themselves,  falling  between 
the  very  strong  acid-forming  elements  and  the  equally  strong  base- 
forming  ones.  It  was  therefore  very  interesting  that  they  should  have 
no  chemical  activity,  but  should  serve  as  a  sort  of  bridge  between  the 
elements.  Rather  than  renumber  the  old  columns,  thus  creating  some 
confusion,  the  new  one  was  added  as  column  O,  this  symbol  having  a 
certain  fitness,  since  the  elements  in  it  have  no  chemical  activity. 

Two  families  in  a  group.  A  study  of  the  elements  comprised  within 
a  group  will  disclose  the  fact  that  each  group  (excepting  Group  O) 
falls  naturally  into  two  families.  The  elements  in  the  odd-numbered 
rows,  or  series,  form  one  family,  those  in  the  series  of  even  numbers, 
the  other.  In  the  table  these  are  arranged  under  the  headings 
A  and  B.  The  elements  in  one  family  are  much  more  similar  to  each 
other  than  they  are  to  those  in  the  other  family  in  the  same  group. 
Thus,  magnesium,  zinc,  cadmium,  and  mercury  form  the  one  family 
of  very  similar  elements  in  Group  II,  while  calcium,  strontium,  barium, 
and  radium  form  the  other. 


CLASSIFICATION  OF  THE  ELEMENTS 


237 


ss 

si 

a  S 


0" 


4 


lii 

II  II  II 


238  GENERAL  CHEMISTRY 

The  first  series  are  type  elements.  The  first  row  of  elements  of 
smallest  atomic  weight  do  not  fall  distinctively  into  either  family  in 
their  group,  but  seem  to  combine  the  properties  of  each,  and  so  serve 
as  types  of  the  general  characteristics  of  the  group.  Thus,  glucinum, 
in  Group  II,  has  the  same  valence  as  all  other  elements  in  the  group, 
and  has  the  general  characteristics  of  both  the  calcium  and  the  magne- 
sium family,  but  it  could  not  be  properly  classed  with  either  one.  These 
type  elements  are  therefore  placed  in  the  middle  of  each  column,  and 
not  in  either  family  at  the  two  sides. 

Group  resemblances.  In  any  one  group  the  elements  have  much  the 
same  valence  characteristics.  They  may  therefore  be  expected  to  form 
oxides,  hydroxides,  acids,  and  salts  of  the  same  general  formula.  The 
hydrides  and  oxides  of  maximum  valence  are  indicated  by  the  for- 
mulas at  the  top  of  the  columns.  It  will  be  seen  that  the  valence 
toward  oxygen  increases  from  0  in  Group  O  to  8  in  Group  VIII, 
while  that  toward  hydrogen  reaches  a  maximum  in  Group  IV.  For 
example,  sulfur,  in  the  sixth  column,  has  an  oxygen  valence  of  6,  as 
shown  in  the  oxide  of  the  formula  SO3,  while  its  valence  toward  hydro- 
gen is  2,  as  indicated  in  the  formula  H2S.  It  usually  happens  that,  in 
addition  to  the  oxide  normally  characterizing  a  group,  each  element 
forms  other  oxides  not  forecast  by  the  table.  In  most  cases  about  the 
same  variety  is  exhibited  by  all  the  elements  in  a  group,  though  the 
ones  of  greatest  atomic  weight  are  likely  to  have  the  greatest  variety 
and  to  differ  most  widely  from  the  others.  Thus,  in  the  variety  and 
character  of  its  compounds,  tellurium  differs  more  from  selenium  than 
the  latter  does  from  sulfur. 

Family  resemblances.  The  resemblances  of  the  members  of  a  family 
to  each  other  are  more  marked,  and  extend  not  only  to  the  general 
chemical  characteristics  but  to  physical  properties  as  well.  There  is, 
in  most  cases,  a  regular  variation  in  any  given  property  as  we  pass 
from  the  member  of  smallest  atomic  weight  to  that  of  largest,  so  that 
the  middle  one  is  a  mean  between  the  others.  This  is  true  in  regard  to 
chemical  activity  and  such  properties  as  melting  point,  boiling  point, 
density,  color,  solubility  of  parallel  salts,  and  many  similar  properties. 
These  points  will  be  dwelt  upon  in  the  several  families  as  they  are 
taken  up  in  order. 

Curves  of  physical  constants.  If  we  select  almost  any  property  of 
the  elements,  such  as  melting  point,  density,  atomic  volume  (atomic 
weight  divided  by  density),  compressibility,  or  melting  point  of  chloride, 


CLASSIFICATION  OF  THE  ELEMENTS 


239 


and  plot  the  value  of  this  property  as  ordinate  against  the  value  of 
the  atomic  weight  as  abscissa,  the  periodic  character  of  the  property 
is  very  strikingly  represented.  The  curves  form  a  series  of  crests 
and  hollows,  as  shown  in  Fig.  87,  which  represents  the  atomic 
volumes.  It  will  also  be  noticed  that  members  of  the  same  family 
occupy  analogous  positions  on  the  successive  waves  of  the  curve. 


60 


70 


50 


30 


\Cs 


n/Jf    \ 

%g  fr 

W" 


ATOMIC  WEIGHTS 


0  20  40  60  80  10O  120  140 

FIG.  87 


ISO          ZOO          220       240 


Irregularities  in  the  table.  Notwithstanding  the  many  very  striking 
relations  suggested  by  the  periodic  grouping,  it  contains  a  number  of 
imperfections  and  anomalies. 

In  the  first  place,  the  relations  are  only  qualitative  in  character. 
The1  atomic  weights  themselves  do  not  differ  regularly  from  element 
to  element,  the  difference  in  some  cases  being  very  small,  as  between 
cobalt  and  nickel,  while  in  others  it  is  as  much  as  four  units,  as 
between  calcium  and  scandium.  The  same  irregularity  characterizes 
all  the  relations  represented  in  the  table,  so  that  no  property  of  one 
element  can  be  quantitatively  calculated  from  the  known  magnitude 
of  the  same  property  in  another  element. 

A  most  striking  fact  is  that  there  is  no  place  in  the  table  for  hydro- 
gen, one  of  the  most  important  of  all  the  elements.  It  seems  to  stand 
quite  alone  and  to  have  no  close  relation  to  any  other  element.  It 
certainly  does  not  belong  in  Group  O,  for  it  has  very  pronounced 


240  GENERAL  CHEMISTRY 

chemical  activity.  All  its  physical  characteristics  would  exc.ly.de  it 
from  Group  I,  which  is  composed  of  typical  metals.  It  is  the  only 
gaseous  element  which  forms  positive  ions  in  solution,  and  it  is  only 
in  this  respect  that  it  suggests  metallic  properties.  It  should  be  remem- 
bered, however,  that  it  is  the  lightest  of  all  elements,  and  it  may  well 
be  that,  in  the  nebulae  of  space  and  in  the  atmosphere  of  much  larger 
stars,  like  the  sun,  there  are  other  lighter  gases  which  are  the  analogues 
of  hydrogen.  Indeed,  there  is  some  spectroscopic  evidence  that  this 
is  the  case. 

In  a  few  cases  the  order  in  which  two  elements  are  placed  in  the 
table  is  not  in  keeping  with  the  magnitude  of  their  respective  atomic 
weights.  Thus,  argon  and  potassium  are  reversed,  in  order  that  they 
may  fall  in  the  columns  where  they  undoubtedly  belong.  The  same 
is  true  of  tellurium  and  iodine,  and  of  nickel  and  cobalt.  While  it  is 
possible  that  the  values  of  these  weights  are  imperfectly  determined, 
they  have  been  the  subject  of  so  many  researches  that  it  does  not 
seem  probable  that  they  will  be  greatly  changed. 

The  last  column  in  the  table  has  little  meaning,  though  there  is  a 
certain  regularity  in  the  periodic  reappearance  of  sets  of  three  closely 
similar  elements  at  the  end  of  alternate  series,  and  some  regularities 
can  be  traced  between  corresponding  members  of 'these  sets. 

The  periodic  arrangement  places  sodium,  copper,  silver,  and  gold 
together  in  one  family,  whereas  there  is  not  much  resemblance  in 
the  chemistry  of  any  of  these  elements.  They  are  also  placed  in  the 
family  of  unit  valence,  although  both  copper  and  gold  usually  have 
a  higher  valence. 

One  of  the  most  irregular  features  of  the  grouping  is  in  connection 
with  a  series  of  elements  known  collectively  as  the  rare  earths.  It 
will  be  noticed  that,  in  the  third  group  and  the  seventh  series,  instead 
of  one  element  there  is  indicated  La  139.0  —  Lu  174.0,  with  a  gap 
of  35  units  between  their  atomic  weights.  Between  these  two  limits 
there  are  at  least  fifteen  known  elements,  or  enough  to  reach  more 
than  across  the  table.  Yet  each  of  these  elements  has  a  valence  of  3, 
and  they  are  so  very  similar  that  they  can  hardly  be  separated  at  all. 
Evidently  they  all  belong  in  the  same  family,  and  should  not  be  dis- 
tributed throughout  the  other  families,  to  which  they  have  little  or  -no 
resemblance.  Cerium  alone  possesses  properties  which  warrant  plac- 
ing it  in  a  different  group  (the  fourth).  It  would  seem  that  through- 
out these  fifteen  elements  the  normal  increase  in  atomic  weights  is 


CLASSIFICATION  OF  THE  ELEMENTS  241 

attended  with  no  appreciable  change  in  chemical  character.  The 
regular  order  is  then  resumed  when  tantalum  is  reached. 

Meaning  of  imperfections.  That  some  very  fundamental  relations 
are  represented  in  the  periodic  table  cannot  be  doubted.  It  can  hardly 
be  mere  chance  which,  on  arranging  the  elements  in  accordance  with 
their  atomic  weights,  brings  into  the  same  group  those  which  have  long 
been  recognized  as  intimately  related.  At  the  same  time  the  unques- 
tioned irregularities  in  the  table  suggest  that  the  present  arrange- 
ment is  far  from  perfect,  and  that  the  real  relationships  are  much  more 
intricate  than  that  which  underlies  the  periodic  grouping. 

Many  attempts  have  been  made  to  improve  upon  the  plan  of  arrange- 
ments, and  some  of  the  proposed  plans  are  very  interesting  and  sug- 
gestive, especially  those  in  which  the  elements  are  arranged  on  a  spiral 
curve,  or  a  helix,  in  accordance  with  a  logarithmic  law.  It  would  lead 
too  far  to  describe  such  plans,  and  a  larger  work  must  be  consulted 
for  further  information.1 

Value  of  the  periodic  law.  While  admitting  that  the  table  contains 
many  imperfections  and  will  doubtless  undergo  many  modifications, 
it  is  nevertheless  of  very  great  service  to  us  in  many  ways. 

It  is  a  constant  check  on  experimental  work.  If  a  given  piece  of 
work  results  in  values  which  are  not  in  harmony  with  the  table,  atten- 
tion is  at  once  directed  to  the  value,  and  it  is  very  carefully  verified. 
Thus,  it  was  at  one  time  announced  that  the  density  of  pure  caesium 
was  2.4.  This  value  was  higher  than  would  be  expected  from  the 
values  for  potassium  and  rubidium,  and  a  careful  redetermination 
gave  the  value  1.87. 

The  service  of  the  table  can  best  be  seen  by  going  back  to  the  time 
before  the  periodic  law  was  formulated.  There  was  then  no  guide  at 
all  as  to  probable  values,  and  one  value  for  a  physical  or  chemical 
property  was  as  reasonable  as  any  other.  When  Mendeleeff  first 
arranged  the  table,  a  number  of  elements  were  plainly  out  of  place, 
and  many  of  their  physical  constants  were  out  of  keeping  with  the 
position  of  the  element.  In  almost  every  case  a  careful  study  showed 
that  the  atomic  weight  was  wrong  or  the  property  erroneously  deter- 
mined, the  revised  values  coming  into  harmony  with  the  table. 

The  table  has  also  been  of  value  in  forecasting  the  discovery  of 
new  elements,  together  with  their  probable  character.  In  one  or  two 
cases  it  has  even  been  possible  to  predict  in  what  kind  of  mineral  they 

i  See  Venable,  The  Periodic  Law. 


242  GENERAL  CHEMISTRY 

would  probably  be  found.  In  arranging  the  table  it  was  found  neces- 
sary to  leave  certain  positions  blank,  in  order  that  the  next  element 
should  fall  in  the  column  where  it  evidently  belonged.  These  blanks 
were  obviously  the  positions  of  undiscovered  elements,  and  from  their 
position  and  the  character  of  the  elements  on  each  side  of  them  their 
properties  could  be  in  a  measure  predicted.  In  this  way  Mendeleeff 
was  able  to  predict  the  properties  of  scandium,  gallium,  and  germanium 
before  they  were  discovered.  Some  blank  spaces  still  remain,  and  two 
have  been  filled  in  recent  years  by  the  discovery  of  radium  and  niton. 

On  the  other  hand,  the  table  has  been  of  much  service  in  ruling  out 
a  number  of  elements  for  which  there  was  no  place.  Many  of  these 
have  been  described  from  time  to  time,  and  as  soon  as  it  was  seen  that 
there  was  no  place  for  them,  they  became  the  center  of  much  interest. 
This  resulted  in  more  extended  study,  and  experiment  soon  showed 
that  the  supposed  element  was  not  entitled  to  such  a  rating. 

The  very  fact  that  the  table  is  imperfect  has  stimulated  a  vast 
amount  of  careful  work,  to  secure  more  accurate  results  on  the  con- 
stants of  the  elements,  so  that  a  more  perfect  generalization  may 
be  reached.  At  the  present  time  the  atomic  weights  are  being  re- 
determined  with  a  care  never  before  devoted  to  them,  and  constants 
of  other  kinds,  such  as  melting  and  boiling  points,  densities,  and 
compression  coefficients,  are  also  undergoing  revision. 

Finally,  from  the  standpoint  of  convenience  the  table  is  of  much 
service  in  assisting  the  memory.  When  the  general  character  of  the 
relations  between  the  elements  is  understood,  one  is  reasonably  safe 
in  assuming  that  in  all  probability  a  given  element  will  have  about 
the  properties  suggested  by  its  position.  Such  an  assumption  is 
always  open  to  doubt,  but  if  merely  general  points  are  involved, 
and  not  quantitative  relations,  the  assumption  is  usually  borne  out 
by  the  facts. 


CHAPTER  XIX 

THE  CHLORINE  FAMILY 


ATOMIC 
WEIGHT 

MELTING 
POINT 

BOILING 
POINT 

COLOR  AND  STATE 

Fluorine  (F) 

19.00 

-223° 

-187° 

Pale  yellowish  gas 

Chlorine  (Cl)     .     .     . 

35.46 

-102° 

-  33.6° 

Greenish-yellow  gas 

Bromine  (Br)     . 

79.92 

-7° 

63° 

Red  liquid 

Iodine  (I)      .... 

126.92 

113° 

184.4° 

Purplish-black  solid 

Characteristics  of  the  family.  The  four  elements  named  in  the  table 
form  a  strongly  marked  family  and  illustrate  very  clearly  the  way  in 
which  the  members  of  a  periodic  family  resemble  each  other,  as  well  as 
the  character  of  the  differences  which  we  may  expect  to  find  among 
the  several  individuals.  Before  taking  up  a  discussion  of  each  of  these 
elements,  it  is  desirable  to  discuss  the  group  as  a  whole,  pointing 
out  the  relation  which  the  properties  of  the  individual  members  bear 
to  each  other,  as  well  as  the  character  of  some  of  their  compounds. 

1.  Occurrence.   Because  of  their  great  activity  these  elements  do  not 
occur  in  the  free  state  in  nature.    Their  compounds,  however,  are  very 
abundant,  those  of  chlorine,  bromine,  and  iodine  being  found  especially 
in  sea  water.    The  most  abundant  of  these  is  sodium  chloride,  or  com- 
mon salt.    Because  the  other  elements  of  the  family  fornr  compounds 
resembling  common  salt,  they  are  often  termed  the  halogens,  a  word 
meaning  "  producers  of  salt." 

2.  Properties.    In  connection  with  the  periodic  law  it^was  pointed 
out  that  the  elements  constituting  a  family  exhibit  a  more  or  less 
regular  gradation  of  properties.    In  the  case  of  the  elements  of  the 
chlorine  family  this  characteristic  is  readily  observed  by  reference  to 
the  table.    It  will  be  seen  that  the  melting  points  and  boiling  points 
of  these  elements  are  in  the  same  order  as  then-  atomic  weights.    A 
somewhat  similar  gradation  is  noted  in  the  color  of  the  elements,  as 
well  as  in  their  physical  states.    Thus,  while  both  fluorine  and  chlorine 
are  gases,  chlorine  is  much  more  readily  condensed  and  has  a  much 
deeper  color  than  fluorine ;  bromine  is  a  brownish-red  liquid,  while 

243 


244  GENERAL  CHEMISTEY 

iodine  is  a  purplish-black  solid.  A  similar  gradation  of  properties  is 
noted  in  their  chemical  conduct.  For  example,  the  affinity  of  these 
elements  for  hydrogen  and  metals  under  the  same  conditions  is  in  the 
reverse  order  of  their  atomic  weights,  being  greatest  in  the  case  of 
fluorine  and  least  in  the  case  of  iodine.  It  follows  that  the  compounds 
of  fluorine  with  hydrogen  and  metals  are  in  a  general  way  the  most 
stable,  while  the  corresponding  compounds  of  iodine  are  the  least  stable. 

The  affinity  of  the  elements  of  the  family  for  oxygen  is  very  slight. 
The  most  stable  oxide  of  the  group  is  iodine  pentoxide  (I2O5),  and  even 
this  is  decomposed  with  ease.  Chlorine  forms  three  oxides  (C12O,  C1O2, 
and  C12O7),  all  of  which  are  very  unstable.  No  oxides  of  fluorine 
or  bromine  are  known  with  certainty. 

3.  Compounds  with  hydrogen.  Hydrogen  combines  with  each  of  the 
elements  of  the  family  to  form  the  following  important  compounds : 

•     Hydrogen  fluoride  (H2F2)  :  a  colorless  liquid  boiling  at  19.4°. 

Hydrogen  chloride  (HC1) :  a  colorless  gas  condensing  to  a  liquid  at  —  83.1°. 
Hydrogen  bromide  (HBr) :  a  colorless  gas  condensing  to  a  liquid  at  —  73°. 
Hydrogen  iodide  (HI)  :  a  colorless  gas  condensing  to  a  liquid  at  —  34.1°. 

These  compounds,  in  the  complete  absence  of  water,  are  rather  inac- 
tive and  have  neither  acid  nor  basic  properties.  They  dissolve  in  water, 
however,  forming  solutions  that  are  acid  in  character.  These  solutions 
are  known  respectively  as  hydrofluoric,  hydrochloric,  hydrobromic,  and 
hydriodic  acids.  Hydrofluoric  acid  is  rather  weak,  but  the  three  re- 
maining ones  are  among  the  strongest  known.  The  salts  of  these  acids 
are  known  respectively  as  fluorides,  chlorides,  bromides,  and  iodides. 

In  their  compounds  with  hydrogen  and  the  metals,  the  halogens 
are  univalent.  In  their  oxides,  chlorine  and  iodine  have  a  much 
higher  maximum  valence,  chlorine  being  apparently  heptavalent  in  the 
oxide  C12O7.  It  may  be  added  that  fluorine  is  not  so  closely  related 
in  its  properties  to  chlorine,  bromine,  and  iodine  as  the  latter  elements 
are  to  each  other,  as  is  illustrated  by  a  comparison  of  the  properties 
of  their  acids. 

FLUORINE 

History  and  occurrence.  The  most  common  forms  in  which  fluorine 
occurs  in  nature  are  the  minerals  known  respectively  as  fluor  spar, 
cryolite,  and  fluorapatite.  Fluor  spar  is  calcium  fluoride  (CaF2).  It  is 
widely  distributed  and  is  found  in  large  quantities,  especially  in  Illinois. 
Cryolite  is  a  fluoride  of  sodium  and  aluminium  (Na3AlF6)  and  is 
found  in  Greenland  and  Iceland.  Fluorapatite  is  the  most  abundant 


THE  CHLORINE  FAMILY 


245 


of  the  three  and  consists  apparently  of  calcium  fluoride  combined  with 
calcium  phosphate,  as  represented  in  the  formula  3  Ca3(PO4)2  •  CaF  . 
Traces  of  compounds  of  fluorine  are  also  found  in  sea  water,  in  many 
minerals,  in  bones,  and  in  the  enamel  of  the  teeth. 

While  the  compounds  of  fluorine  have  long  been  known,  all  efforts 
to  liberate  the  element  failed  until  the  year  1886,  when  the  French 
chemist  Moissan  finally  succeeded  in  isolating  it  and  made  an  ex- 
tended study  of  its  properties. 

Preparation.  Because  of  its  great  activity,  fluorine  cannot  be  pre- 
pared by  any  of  the  general  methods  which  serve  for  the  preparation 
of  the  other  members  of  the  chlorine  family.  Moissan  finally  isolated 
it  by  the  electrolysis  of  hydrogen  fluoride.  Pure  hydrogen  fluoride, 
like  pure  water,  is  not  an  elec- 
trolyte. As  in  the  case  of  water, 
it  may  be  made  an  electrolyte  by 
dissolving  in  it  an  appropriate 
compound.  For  this  purpose 
Moissan  used  potassium  hydro- 
gen fluoride  (KHF2),  a  solid  that 
is  readily  obtained  in  a  pure  con- 
dition. Since. hydrogen  fluoride 
boils  at  19.4°,  the  operation  must 
be  carried  out  at  a  low  tempera-  • 
ture,  to  prevent  the  liquid  from 
vaporizing.  On  passing  a  current 
of  electricity  through  the  solu- 
tion, hydrogen  is  evolved  at  the 
cathode  and  fluorine  at  the  anode. 

Moissan's  method.  The  solution  of 
potassium  hydrogen  fluoride  in  hydro- 
gen fluoride  was  introduced  into  a  U-shaped  tube  A  (Fig.  88)  made  of  platinum. 
The  tube  was  fitted  with. calcium  fluoride  stoppers  B,  B,  through  which  passed 
wires  attached  to  electrodes  C,  C,  made  of  an  alloy  of  platinum  and  iridium,  which 
is  very  resistant  to  the  action  of  reagents.  The  tube  containing  the  solution  was 
placed  in  a  vessel  D  filled  with  a  low-boiling  liquid  which,  by  its  rapid  evapora- 
tion, reduced  the  temperature  sufficiently  to  prevent  the  vaporization  of  the  hydro- 
gen fluoride.  On  electrolyzing  the  solution,  fluorine  was  evolved  at  the  anode  and 
escaped  through  the  tube  E,  while  the  hydrogen  evolved  at  the  cathode  escaped 
through  the  tube  F.  Moissan  found  later  that  copper  is  but  slightly  attacked  by 
fluorine,  and  that  a  copper  tube  could  therefore  be  used  in  place  of  the  more 
expensive  platinum  tube. 


FIG.  88 


246  GENERAL  CHEMISTRY 

Properties.  Fluorine  is  a  gas,  slightly  yellow  in  color.  It  is  1.3 
times  as  heavy  as  air.  It  can  be  obtained  in  the  form  of  a  yellow 
liquid  which  boils  at  —  187°  and  solidifies  at  —  223°.  Chemically  it 
is  one  of  the  most  active  of  all  elements.  Most  of  the  metals  and 
many  of  the  nonmetals,  when  brought  in  contact  with  fluorine,  com- 
bine with  it  so  rapidly  as  to  produce  -light.  It  unites  with  hydrogen 
with  explosive  violence  and  readily  abstracts  it  from  its  compounds. 
For  example,  it  decomposes  water  violently,  forming  hydrogen  fluoride 
and  oxygen:  2F2  +  2H2O  =  2H2F2  +  O2  .  '_ 

From  10  to  14  per  cent  of  the  oxygen  liberated  is  at  the  same  time 
converted  into  ozone.  It  liberates  all  the  other  members  of  the  chlo- 
rine family  from  their  compounds  with  hydrogen  and  the  metals.  It 
does  not  combine  with  oxygen,  however,  and  gold,  platinum,  and  copper 
are  but  slightly  attacked  by  it. 

Hydrogen  fluoride  (H2F2).  Pure  hydrogen  fluoride  is  best  prepared 
by  heating  anhydrous  potassium  hydrogen  fluoride: 

2  KHF2  =  2  KF  +  H2F2 

The  process  must  be  carried  out  in  platinum  vessels.  Hydrogen  fluo- 
ride can  also  be  prepared  by  the  action  of  sulfuric  acid  on  the  fluorides 
of  the  metals.  Calcium  fluoride,  being  the  cheapest,  is  always  used  : 


The  operation  is  carried  out  in  vessels  of  platinum  or  lead.  The 
hydrogen  fluoride  so  obtained  is  not  entirely  anhydrous.  This  is  the 
method  commonly  used  in  the  preparation  of  its  aqueous  solutions, 
in  which  form  it  is  generally  used. 

Properties.  Hydrogen  fluoride  is  a  colorless  liquid  that  boils  at  19.4°. 
Its  vapor  is  lighter  than  air.  At  low  temperatures  it  forms  a  white 
solid  melting  at  —92.3°.  An  aqueous  solution  containing  35  per  cent 
of  the  compound  has  a  constant  boiling  point,  namely,  120°,  and  distills 
with  unchanged  concentration  (p.  127).  As  indicated  by  the  formation 
of  such  acid  salts  as  KHF2,  its  formula  is  generally  regarded  as  H2F2, 
although  at  different  temperatures  it  may  exist  in  any  of  the  forms 
indicated  by  the  formulas  _ 

HF>      H2F2>      H3F3 

Pure  hydrogen  fluoride  is  a  nonconductor  of  electricity  and  is  neu 
tral  in  reaction.    When  dissolved  in  water,  it  ionizes  as  follows: 

H2F2:«=>:H+,  H+  +  F-,  F- 


THE  CHLORINE  FAMILY  24T 

The  resulting  solution  is  acid  in  character,  owing  to  the  presence  of 
hydrogen  ions,  and  is  known  as  hydrofluoric  acid.  Ordinary  commer- 
cial hydrofluoric  acid  contains  about  50  per  cent  of  hydrogen  fluoride. 
Since  this  solution  readily  attacks  glass,  it  is  kept  in  bottles  made  of 
wax,  ceresin  being  generally  used. 

Hydrofluoric  acid.  Hydrofluoric  acid  is  a  rather  weak  acid  and  pos- 
sesses all  the  characteristics  of  such  a  compound.  It  acts  upon  some  of 
the  metals,  as  well  as  upon  their  oxides  and  hydroxides,  forming  the 
corresponding  salts.  A  distinguishing  property  of  hydrofluoric  acid  is 
its  action  upon  glass.  Ordinary  glass  consists  principally  of  sodium 
silicate  (Na2SiO3)  and  calcium  silicate  (CaSiO3).  Hydrofluoric  acid 
reacts  with  these  silicates,  forming  fluorides  of  sodium  and  calcium, 
while  the  silicon  present  also  combines  with  fluorine  to  form  tetra- 
fluoride  of  silicon,  a  gaseous  compound  having  the  formula  SiF4 : 

Na2Si03  +  3  H2F2  =  Na2F2  +  SiF4  +  3  H2O 
CaSi03  +  3  H2F2  =  CaF2  +  SiF4  +  3  H2O 

Advantage  is  taken  of  this  reaction  in  marking  thermometers  and 
flasks,  as  well  as  in  etching  designs  on  glassware. 

The  etching  property  of  the  acid  may  be  shown  in  the  following  way  :  A  piece 
of  glass  is  covered  with  a  thin  coating  of  some  wax,  such  as  paraffin.  To  do  this 
the  wax  is  melted  on  the  glass,  which  is  then  tipped  until  the  melted  wax  com- 
pletely covers  it  with  a  thin  film.  After  the  wax  hardens,  any  desired  design  is 
made  by  cutting  through  the  wax  with  a  fine  metal  point.  The  object  so  pre- 
pared is  then  exposed  to  the  fumes  of  hydrofluoric  acid.  The  acid  attacks  the 
glass  wherever  exposed,  destroying  its  luster  and  leaving  the  design  etched  upon 
its  surface. 

Hydrofluoric  acid  is  exceedingly  corrosive.  A  single  drop  in  con- 
tact with  the  skin  produces  a  very  painful  wound,  slow  to  heal.  Its 
vapor  must  not  be  inhaled. 

Salts  of  hydrofluoric  acid ;  the  fluorides.  The  fluorides  of  most  of 
the  metals  are  known.  They  can  be  prepared  by  the  usual  method  for 
preparing  salts,  namely,  by  the  action  of  the  acid  upon  the  metals 
directly  or  upon  their  oxides  or  hydroxides.  The  most  import-ant  of 
these  fluorides  is  the  well-known  calcium  fluoride,  or  fluor  spar,  as  it 
is  termed  when  found  in  nature.  A  number  of  the  metals  form  acid 
salts  such  as  KHF2. 


248  GENERAL  CHEMISTRY 

CHLORINE 

History  and  occurrence.  Scheele,  who  was  the  first  to  obtain  oxygen 
in  the  pure  state,  was  likewise  the  first  to  isolate  chlorine  (1774).  He 
obtained  the  element  by  the  action  of  hydrochloric  acid  upon  manga- 
nese dioxide,  a  method  of  preparation  which  is  still  used.  The  element 
was  regarded,  however,  as  a  compound  of  hydrochloric  acid  with  oxy- 
gen until  the  English  chemist  Davy,  in  1810,  proved  its  elementary 
character. 

Because  of  its  color  Davy  named  the  element  chlorine,  from  the 
Greek  word  meaning  "  greenish  yellow." 

The  most  abundant  compound  of  chlorine  is  sodium  chloride,  or 
common  salt.  This  compound  is  found  in  sea  waters  and  in  large 
deposits  in  various  parts  of  the  earth.  Chlorine  also  occurs  in  nature 
in  combination  with  potassium,  magnesium,  calcium,  and,  to  a  limited 
extent,  with  some  of  the  other  metals.  Sodium  chloride  is  an  essential 
constituent  of  our  food,  while  the  acid  character  of  the  gastric  juice 
is  due  largely  to  hydrochloric  acid. 

Preparation.  Three  general  methods  are  in  use  for  the  preparation 
of  chlorine.  A  discussion  of  each  of  these  follows : 

1.  Preparation  by  the  decomposition  of  per  chlorides.  When  a  solution 
of  hydrochloric  acid  is  gently  heated  with  manganese  dioxide,  man- 
ganese tetrachloride  forms,  which  is  unstable  and  decomposes  with  the 
evolution  of  free  chlorine,  as  expressed  in  the  following  equations : 

MnO2  +  4  HC1  =  2  H2O  4-  MnCl4 
MnCl4  =  MnCl2  +  C12 

Since  sodium  chloride  and  sulfuric  acid  interact  to  form  hydrochloric 
acid,  it  is  often  found  more  convenient  to  substitute  a  mixture  of  these 
two  compounds  for  the  hydrochloric  acid  in  the  above  process.  The 
reaction  is  then  expressed  in  the  following  equation : 

2  NaCl  +  3  H2S04  +  MnO2  -  2  NaHSO4  +  2  H2O  +  MnSO4  +  C12 

This  general  method  serves  as  a  convenient  one  for  preparing  chlorine 
in  the  laboratory.  It  may  be  carried  out  as  follows : 

Manganese  dioxide  and  hydrochloric  acid  are  placed  in  the  flask  A  (Fig.  89)  and 
thoroughly  mixed.  A  gentle  heat  is  then  applied  to  the  flask.  Chlorine  is  evolved 
and,  escaping  through  the  tube  B,  bubbles  through  the  water  in  bottle  C  (which 
removes  any  hydrogen  chloride  carried  over  with  it)  and  finally  through  some 
sulfuric  acid  in  bottle  D  (which  removes  any  moisture  present).  Since  the  gas  is 


THE*  CHLORINE  FAMILY 


249 


fairly  soluble  in  water,  it  is  collected  by  displacing  the  air  in  bottles  or  cylinders  E. 
Because  of  the  poisonous  character  of  chlorine,  the  preparation  must  be  carried 
on  in  a  well-ventilated  hood. 

2.  Action  of  oxidizing  agents  upon  hydrogen  chloride.  Under  suitable 
conditions,  oxygen  acts  upon  hydrogen  chloride,  liberating  chlorine  in 
accordance  with  the  following  equation  : 

0=2H0  +  2C1 


Either  free  oxygen  itself  or  some  compound  which  readily  evolves 

oxygen  may  be  used.    The  compounds  most  frequently  employed  are 

potassium  dichromate  and  potassium  permanganate.   The  speed  of  the 

reaction  between  free 

oxygen   and   hydro- 

gen chloride  is  very 

slow.    By  the  use  of 

a  suitable   catalytic 

agent,   such  as    the 

chloride  or  sulf  ate  of 

copper,  the  speed  may 

be  greatly  increased, 

free   chlorine    being 

liberated  in  accord- 

ance with  the  above 

equation.    The  cata- 

lyzer is  prepared  by 

saturating  some  porous  material,  such  as  pieces  of  bricks,  with  a  solu- 

tion of  the  copper  compound.   This  is  then  placed  in  a  tube  and  the 

mixture  of  hydrogen  chloride  and  air  passed  through  the  tube,  which 

is  heated  to  about  375°.    This  process,  which  is  known  as  Deacon's 

process,  is  used  to  a  limited  extent  in  England  for  the  preparation  of 

chlorine  on  a  commercial  scale,  but  it  has  never  come  into  use  in  the 

United  States.    The  chlorine  so  obtained  is  not  pure,  being  mixed 

with  the  nitrogen  from  the  air  and  with  excess  of  hydrogen  chloride, 

but  is  suitable  for  many  purposes. 

In  the  laboratory  it  is  much  more  convenient  to  use  hydrochloric  acid  in  place 
of  hydrogen  chloride,  and  either  potassium  permanganate  (KMnO4)  or  potassium 
dichromate  (K2Cr2O7)  as  an  oxidizing  agent  in  place  of  free  oxygen.  Each  of 
these  compounds  reacts  with  hydrochloric  acid,  liberating  free  chlorine.  The 
complete  reactions  are  somewhat  complicated  and,  will  be  discussed  in  a  later 
chapter. 


FIG.  89 


250 


GENERAL   CHEMISTRY 


FIG.  90 


To  obtain  chlorine  in  the  laboratory  by  this  method  the  potassium  perman- 
ganate is  placed  in  a  flask  A  (Fig.  90),  and  a  mixture  of  equal  volumes  of 
concentrated  hydrochloric  acid  and  water  is  added,  drop  by  drop,  from  a  separa- 
tory  funnel  B.  The  reaction  takes  place  at  once,  and  the  evolved  chlorine  may 

be  collected  by  displacement  of  air.    If  a  high 
degree    of   purity  is   desired,  the   gas  should  be 
B  passed  through  sulfuric  acid. 

3.  Electrolytic  methods.  Chlorine  is  read- 
ily obtained  by  the  electrolysis  of  a  solution 
of  sodium  chloride  (p.  143).  This  is  the 
method  most  generally  used  for  the  prepa- 
ration of  chlorine  on  a  large  scale.  At  the 
present  time  all  the  chlorine  prepared  for 
commercial  purposes  in  the  United  States, 
and  much  of  that  in  Europe,  is  obtained  in 
this  way. 

The  electrolytic  method  possesses  the  following  advantages :  (1)  sodium 
chloride  is  cheap;  (2)  in  addition  to  chlorine,  sodium  hydroxide,  for  which 
there  is  great  demand,  is  formed  in  the  process.  The  chief  item  of  cost  is  the 
generation  of  the  electrical  energy,  so  that  naturally  the  factories  for  the  pro- 
duction of  chlorine  are  located  at  points  where  water  power  can  be  used  to 
advantage,  as  at  Niagara  Falls.  The  only  reason  why  the  electrical  method  is 
not  universally  used  is  due  to  the  fact  that  in  parts  of  Europe  large  quantities 
of  hydrochloric  acid  are  obtained  in  the  manufacture  of  sodium  carbonate  by 
the  Leblanc  process,  and  the  most  economical  use  found  for  this  acid  is  in  the 
preparation  of  chlorine. 

By  far  the  largest  amount  of  the  chlorine  prepared  commercially 
is  used  in  the  preparation  of  bleaching  powder,  which  will  be  dis- 
cussed in  the  following  chapter.  Free  chlorine,  condensed  in  strong 
iron  cylinders,  is  now  an  article  of  commerce. 

Properties.  Chlorine  is  a  greenish-yellow  gas  and  possesses  a 
peculiar  suffocating  odor.  It  is  2.49  times  as  heavy  as  air,  and  under 
ordinary  conditions  1  volume  of  water  dissolves  about  2.5  volumes 
of  the  gas.  At  ordinary  temperatures  (18°)  it  is  liquefied  by  a 
pressure  of  16.5  atmospheres.  Since  the  commercial  chlorine  stored 
in  iron  cylinders  is  subjected  to  a  much  greater  pressure  than  this, 
it  is  evident  that  in  these  cylinders  it  is  in  a  liquid  state.  Liquid 
chlorine  is  yellowish  in  color,  boils  at  —33.6°,  and  solidifies  at 
—102°.  When  inhaled  in  small  quantities,  chlorine  produces  the 
symptoms  of  a  hard  cold,  and  in  larger  quantities  may  have  serious 
and  even  fatal  effects. 


THE  CHLOKINE  FAMILY  251 

Chemical  conduct.  At  ordinary  temperatures  chlorine  is  more  active 
chemically  than  any  of  the  elements  we  have  so  far  considered,  with 
the  exception  of  fluorine.  Indeed,  it  is  one  of  the  most  active  of  all 
the  elements.  The  compounds  formed  by  its  union  with  any  other 
element  are  called  chlorides. 

1.  Action  upon  metals.  Nearly  all  of  the  metals  combine  directly 
with  chlorine,  especially  when  heated.    A  strip  of  copper  foil,  heated 
to  redness  and  immediately  dropped  into  chlorine,  burns  with  incan- 
descence.   Sodium  burns  brilliantly  when  heated  strongly  with  moist 
chlorine.    Gold  and  silver  are  quickly  tarnished  by  the  gas,  and  even 
platinum  is  readily  attacked  by  it. 

2.  Action  upon  nonmetals.  Chlorine  has  likewise  a  strong  affinity  for 
most  of  the  nonmetals.    Thus,  phosphorus  and  sulfur  burn  in  a  cur- 
rent of  the  gas,  while  antimony  and  arsenic,  in  the  form  of  a  fine 
powder,  at  once  burst  into  flame  when  brought  in  contact  with  it, 
forming  in  each  case  the  chloride  of  the  element. 

3.  Action  upon  hydrogen.    Chlorine  unites  readily  with  hydrogen, 
forming  hydrogen  chloride.    A  jet  of  hydrogen  burning  in  the  air 
continues  to  burn  when  introduced  into-  a  jar  of  chlorine,  giving  a 
somewhat   luminous  flame.    A  mixture  of  the  two  gases  explodes 
violently  either  when  heated  or  when  exposed  to  bright  sunlight. 

4.  Action  upon  compounds  of  hydrogen.   Not  only  will  chlorine  com- 
bine directly  with  free  hydrogen,  but  it  will  remove  the  element  from 
some  of  its  compounds.  Thus,  when  chlorine  is  passed  into  an  aqueous 
solution  of  hydrogen  sulfide,  sulfur  is  precipitated  and  hydrochloric 
acid  formed,  as  expressed  in  the  following  equation : 

H2S  +  C12  =  2  HC1  +  S 
With  ammonia  the  action  is  similar : 


Under  certain  conditions  the  nitrogen  evolved  combines  with  chlorine 
to  form  a  very  explosive,  oily  liquid  known  as  nitrogen  trichloride. 
The  strong  affinity  of  chlorine  for  hydrogen  is  very  strikingly  shown 
by  its  action  upon  turpentine.  This  latter  substance  is  made  up  of 
compounds  containing  carbon  and  hydrogen.  When  a  strip  of  paper 
moistened  with  warm  turpentine  is  placed  in  a  jar  of  chlorine,  the 
hydrogen  and  chlorine  unite,  with  evolution  of  light,  forming  hydro- 
gen chloride,  while  a  black  deposit  of  carbon  remains. 


252  GENERAL  CHEMISTRY 

5.  Action  upon  water.  The  liquid  resulting  from  passing  chlorine 
into  water  is  generally  regarded  simply  as  a  solution  of  chlorine  in 
water  and  is  called  chlorine  water.   It  is  probable,  however,  that  chlorine 
reacts  with  the  water  to  form  a  mixture  of  hydrochloric  acid  (HC1) 
and  hypochlorous  acid  (HC1O),  until  the  equilibrium  expressed  in 
the  following  equation  results : 

C12  +  H20  +=±  HC1  +  HC10  (1) 

Hypochlorous  acid  is  unstable,  however,  and  decomposes,  slowly  in  the 
dark  but  rapidly  in  the  sunlight,  into  hydrochloric  acid  and  oxygen : 

2  HC1O  =  2  HC1  +  O2  (2) 

This  removal  of  the  hypochlorous  acid  through  decomposition  disturbs 
the  equilibrium  expressed  in  equation  (1)  so  that  the  interaction  of 
the  chlorine  and  water  continues  as  long  as  any  free  chlorine  is  left. 
There  finally  results  a  dilute  solution  of  hydrochloric  acid,  as  is 
shown  by  combining  equations  (1)  and  (2)  in 
the  usual  way.  The  resulting  equation  is 

|  2C12  +  2H20  =  4HC1  +  02 

The  effect  of  sunlight  in  increasing  the  action  of 
chlorine  upon  water  may  be  shown  in  the  following 
way :    If  a  long  tube  of  rather  large  diameter  is  filled 
with  a  saturated  solution  of  chlorine  in  water  and 
inverted  in  a  vessel  of  the  same  solution  (as  shown 
in  Fig.  91),  and  the  apparatus  is  placed  in  bright  sun- 
.    light,  bubbles  of  gas  will  soon  be  seen  to  rise  through 
—     the  solution  and  collect  in  the  tube.    An  examination 
FIG.  91  of  this  gas  will  show  that  it  is  oxygen. 

The  decomposition  of  water  through  the  action  of  chlorine  is  also 
greatly  increased  in  the  presence  of  some  substance  which  combines 
with  the  oxygen  as  fast  as  it  is  set  free.  Consequently,  a  solution  of 
chlorine  in  water  is  a  good  oxidizing  agent,  and,  indeed,  it  is  often 
used  as  such. 

6.  Formation  of  hydrates.    When  chlorine  is  passed  into  water  and 
the  solution  is  cooled  to  a  point  just  above  freezing,  a  crystalline  hydrate 
separates,  which  has  the  composition  C12  •  8  H2O.   As  the  temperature 
rises,  the  hydrate  gradually  dissociates  into  its  constituents.   It  is  inter- 
esting to  note  that  it  was  from  this  hydrate  that  Faraday,  in  1823,  first 
obtained  chlorine  in  a  liquid  state,  using  the  form  of  apparatus  shown 
in  Fig.  37  (p.  77). 


THE  CHLORINE  FAMILY  253 

7.  Action  upon  color  substances ;  bleaching  action.    Chlorine  possesses 
a  powerful  bleaching  action.  Strips  of  highly  colored  cloth,  when  moist- 
ened with  water  and  placed  in  jars  of  chlorine,  rapidly  lose  their  color. 
The  presence  of  water  is  essential  to  the  change,  as  may  be  shown  by 
placing  strips  of  the  dry  cloth  in  chlorine  from  which  the  moisture  has 
been  removed  by  bubbling  it  through  sulfuric  acid  (Fig.  89).    Under 
these  conditions  the  color  of  the  cloth  remains  unchanged.   It  is  prob- 
able that  the  bleaching  action  of  chlorine  consists  first  in  its  reaction 
with  water  to  form  hypochlorous  acid.    This  acid  then  decomposes, 
the  resulting  oxygen  reacting  with  the  color  substance  of  the  cloth 
to  form  colorless  compounds.    It  is  evident,  therefore,  that  chlorine 
will  bleach  only  those  materials  the  coloring  matters  of  which  are 
changed  by  its  action  into  colorless  compounds.    It  has  no  bleaching 
action  on  such  color  substances  as  carbon,  and  hence  does  not  affect 
printers'  ink  made  from  carbon.    It  cannot  be  used  for  bleaching 
certain  substances,  like  silk  and  straw,  since  it  injures  the  fabric. 

8.  Action  as  a  germicide.    Chlorine  has  marked  germicidal  properties, 
and  the  free  element,  as  well  as  the  compounds  from  which  it  is  easily 
liberated,  are  used  as  disinfectants. 

Uses  of  chlorine.  As  has  been  stated  above,  chlorine  is  an  excellent 
germicide  and  bleaching  agent,  and  large  quantities  of  the  element  are 
used  for  these  purposes.  The  various  kinds  of  fabrics  woven  from 
vegetable  fibers,  such  as  flax  and  cotton,  are  always  more  or  less  colored 
by  the  presence  of  natural  coloring  matter.  Hence,  if  a  white  fabric 
is  desired,  bleaching  is  necessary.  This  was  formerly  accomplished  by 
exposing  the  fabric  to  the  action  of  the  air  and  sunlight,  but  many  days 
were  required  for  the  completion  of  the  process.  The  same  results  are 
now  obtained  in  a  very  short  time  by  the  use  of  chlorine. 

Chlorine  is  generally  used  commercially  in  the  form  of  bleaching 
powder  (p.  270).  The  chlorine  present  in  this  substance  can  be  liberated 
easily  and  utilized  as  desired.  Increasing  amounts  of  the  free  element 
are  being  used  in  the  preparation  of  certain  of  its  compounds. 

Hydrogen  chloride.  Hydrogen  chloride  may  be  prepared  in  a  num- 
ber of  different  ways,  the  most  important  of  which  are  the  following : 

1.  By  direct  combination  of  its  constituent  elements.  Hydrogen  chloride 
is  formed  by  the  direct  union  of  hydrogen  and  chlorine.  Since  both 
these  elements  are  obtained  in  quantities  in  the  electrolysis  of  solu- 
tions of  sodium  chloride,  this  method  is  used  to  a  limited  extent  in 
the  preparation  of  hydrogen  chloride  on  a  commercial  scale. 


254 


GENERAL   CHEMISTRY 


2.  By  the  action  of  concentrated  sulfuric  acid  upon  chlorides  of  metals. 

Sodium  chloride,  because  of  its  low  cost,  is  always  used.  The  reaction 
is  expressed  by  the  following  equation : 

2  NaCl  +  H2SO4  +=±  Na2SO4  +  2  HC1 

It  will  be  noted  from  the  equation  that  sodium  sulfate  is  likewise 
formed  in  this  process.  Now  the  demand  for  sodium  sulfate  is  very 
great,  large  quantities  of  it  being  used  in  the  preparation  of  sodium 
carbonate,  as  well  as  in  the  manufacture  of  glass.  It  follows  that  in 
the  preparation  of  the  sulfate  large  quantities  of  hydrogen  chloride  are 
produced,  and  indeed  this  is  its  most  important  source.  The  method 

also  serves  as  a  convenient  one 
for  the  preparation  of  hydrogen 
chloride  in  the  laboratory. 

Sodium  chloride  is  placed  in  the 
flask  A  (Fig.  92),  fitted  with  a  funnel 
tube  and  an  exit  tube,  sulfuric  acid  of 
the  proper  concentration  is  added,  and 
the  flask  is  gently  warmed.  Hydrogen 
chloride  is  evolved  and  is  collected  by 
displacement  of  air,  as  in  the  prepara- 
tion of  chlorine.  To  prepare  a  solution 
of  the  gas,  the  end  of  the  exit  tube 
is  fixed  just  above  the  level  of  some 
water  contained  in  a  cylinder  B.  The 
gas,  being  extremely  soluble  in  water, 
is  absorbed  as  fast  as  it  escapes  from 
the  tube.  Care  must  be  taken  not  to  have  the  end  of  the  exit  tube  dip  below  the 
surface  of  the  water,  since  the  solubility  of  the  gas  is  so  great  that  the  water 
would  rush  back  into  the  generating  flask. 

If  the  sulfuric  acid  is  added,  not  to  the  solid  sodium  chloride,  but  to  an  aqueous 
solution  of  the  salt,  there  is  no  very  marked  action.  The  hydrogen  chloride  formed 
is  very  soluble  in  water  and  so  does  not  escape  from  the  solution  unless  heated ; 
hence  a  state  of  equilibrium  is  soon  reached  among  the  four  substances  repre- 
sented in  the  equation. 

3.  By  the  action  of  sodium  chloride  upon  sodium  hydrogen  sulfate.    This 
latter  compound  is  obtained  in  the  manufacture  of  nitric  acid  (p.  176). 
When  heated  with  sodium  chloride,  hydrogen  chloride  is  obtained  in 
accordance  with  the  following  equation  : 

NaHSO4  +  NaCl  ^z±  Na2SO4  +  HC1 

This  method  is  likewise  used  for  the  preparation  of  hydrogen  chloride 
on  a  commercial  scale. 


FIG.  92 


THE  CHLORINE  FAMILY 


255 


FIG.  93 


In  the  commercial  preparation  of  hydrogen  chloride  the  gas  is  ab- 
sorbed in  water,  in  which  it  is  extremely  soluble.  The  resulting  solu- 
tion constitutes  the  ordinary  hydrochloric  acid  of  commerce.  When 
the  materials  are  pure,  the  solution  obtained  is  colorless.  The  com- 
mercial acid,  often  called  muriatic  acid,  is  usually  colored  yellow  by 
impurities. 

The  pure  hydrogen  chloride  can  easily  be  regained  from  this  solution  by  the 
addition  of  sulf  uric  acid,  which  diminishes  the  solubility  of  the  hydrogen  chloride. 
This  serves  as  a  very  convenient  method 
for  obtaining  pure  hydrogen  chloride 
when  a  limited  supply  of  it  is  desired 
in  the  laboratory.  The  concentrated 
solution  is  placed  in  A  (Fig.  93),  and 
the  flask  is  connected  with  the  bottle 
B,  which  contains  sulfuric  acid  for  dry- 
ing the  gas.  This  bottle  is  fitted  with 
stopper  and  tubes,  as  shown  in  the  dia- 
gram. The  glass  tube  leading  from  the 
bottle  B  to  the  bottom  of  the  cylinder 
E  passes  through  a  perforated  card- 
board or  glass  plate  C,  which  rests 
lightly  on  the  top  of  the  cylinder.  Sul- 
furic acid  is  now  added,  drop  by  drop, 
from  the  separatory  funnel  D.  The  hydrogen  chloride  is  at  once  evolved  and, 
after  bubbling  through  the  sulfuric  acid  in  B,  whereby  any  moisture  is  removed, 
is  collected  in  E,  as  shown  in  the  figure. 

By-products.  It  generally  happens  that  in  the  preparation  of  any 
given  substance  other  compounds  are  formed.  These  are  called  by- 
products. Thus,  hydrochloride  acid  is  a  by-product  in  the  manufacture 
of  sodium  sulfate,  just  as  sodium  hydrogen  sulfate  is  a  by-product  in 
the  manufacture  of  nitric  acid.  It  is  evident  that  the  cost  of  the  manu- 
facture of  any  substance  can  be  decreased  to  the  extent  to  which  the 
by-products  can  be  utilized.  The  cost  of  nitric  acid,  for  example,  is 
materially  lessened  by  the  fact  that  sodium  hydrogen  sulfate,  obtained 
along  with  the  nitric  acid,  may  be  used  for  the  preparation  of  sodium 
sulfate  and  hydrochloric  acid,  for  both  of  which  there  is  a  good 
demand.  Indeed,  it  sometimes  happens  that  the  demand  for  the 
by-product  becomes  so  great  that  it  really  comes  to  be  the  main 
product.  The  success  of  a  process  often  depends  upon  the  value 
of  the  by-products  formed. 

Properties  of  hydrogen  chloride.  Hydrogen  chloride,  a  colorless  gas, 
is  1.26  times  as  heavy  as  air.  When  inhaled,  it  has  an  irritating  and 


256 


GENERAL  CHEMISTRY 


suffocating  effect.  At  0°  it  is  condensed  to  the  liquid  state  by  a  pres- 
sure of  28  atmospheres.  The  resulting  liquid  is  colorless,  boils  at 
—  83.1°,  and  solidifies  at  —  113°.  This  liquid  does  not  conduct  elec- 
tricity, has  no  action  upon  metals,  and  in  general  is  very  inactive. 
Hydrogen  chloride  is  very  soluble  in  water,  1  volume  of  the  latter 
under  standard  conditions  dissolving  506  volumes  of  the  gas.  The 
density  of  its  aqueous  solutions  increases  with  the  amount  of  gas 
dissolved,  as  shown  in  the  following  table,  which  gives  the  percent- 
age by  weight  of  hydrogen  chloride  present  in  solutions  of  various 
densities,  the  measurements  being  taken  at  15°. 


PER  CENT 

OF  HC1 

DENSITY 

PEK  CENT 

OF   HC1 

DENSITY 

PER  CENT 

OF  HC1 

DENSITY 

5.69 

1.0284 

20.04 

1.1006 

35.02 

1.1779 

10.17 

1.0507 

25.06 

1.1265 

40.09 

1.2013 

15.22 

1.0761 

30.00 

1.1526 

43.40 

1.2134 

0-0 


Aqueous  solutions  of  hydrogen  chloride  act  like  solutions  of  hydro- 
gen fluoride  upon  distillation  in  that  there  finally  results  a  solution 
of  constant  concentration  and  constant  boiling  point  (p.  126).  In  the 
case  of  the  hydrogen  chloride  this  solution  has  a  concentration  of 
20.24  per  cent  of  the  chloride  and  a  boiling  point  of  110°. 

The  extreme  solubility  of  hydrogen  chloride  in 
water  may  be  shown  as  follows :  A  perfectly  dry 
flask  A  (Fig.  94)  is  filled  with  hydrogen  chloride. 
This  flask  is  connected,  by  means  of  a  glass  tube, 
with  a  similar  flask  B,  which  is  nearly  filled  with 
water,  as  shown  in  the  figure.  The  end  of  the 
tube  opening  into  flask  A  is  drawn  out  to  a  rather 
fine  jet.  By  blowing  into  the  tube  C,  a  few  drops 
of  water  are  forced  into  A .  Some  of  the  hydrogen 
chloride  at  once  dissolves,  thus  diminishing  the 
pressure  inside  the  flask.  The  water  then  flows 
continuously  from  B  into  A,  until  practically  all 
the  hydrogen  chloride  is  absorbed.  It  is  evident 
that  the  connection  must  be  air-tight. 

Composition.  The  composition  of  hydro- 
gen chloride  can  be  determined  by  the  elec- 
trolysis of  its  aqueous  solution.  When 

electrolyzed,  the  hydrogen  of  the  compound  is  evolved  at  the  cathode 
and  the  chlorine  at  the  anode.  A  special  form  of  apparatus  is  required, 
in  order  to  avoid  the  difficulties  arising  from  the  marked  solubility  of 


=  C 


FIG.  94 


THE  CHLORINE  FAMILY  257 

the  chlorine  in  water.  When  the  experiment  is  carried  out,  it  is  found 
that  the  volume  of  the  hydrogen  liberated  is  exactly  equal  to  that  of 
the  chlorine.  Conversely,  it  is  possible  to  show  by  experiment  that 
when  hydrogen  and  chlorine  combine,  they  always  do  so  in  the  ratio  of 
1  volume  of  hydrogen  to  1  volume  of  chlorine ;  moreover,  the  product 
is  always  2  volumes  of  hydrogen  chloride.  These  relations  may  be 
shown  graphically  in  the  following  way : 


+        C12       =       HC1    HC1 

Since  chlorine  is  35.18  times  as  heavy  as  hydrogen,  it  follows  that  1 
part  by  weight  of  hydrogen  combines  with  35.18  parts  by  weight  of 
chlorine  to  form  36.18  parts  by  weight  of  hydrogen  chloride. 

Chemical  conduct  of  hydrochloric  acid.  While  hydrogen  chloride 
itself  has  but  little  chemical  activity,  its  solution  in  water,  namely, 
hydrochloric  acid,  has  marked  chemical  properties  and  constitutes  one 
of  the  most  important  acids.  It  is  relatively  stronger  than  sulfuric 
acid,  being  nearly  equal  to  nitric  acid  in  strength  (p.  155). 

1.  Action  upon  metals  and  upon  their  oxides  and  hydroxides.  Hydro- 
chloric acid  reacts  with  those  metals  that  have  a  higher  electrode 
potential  than  hydrogen,  forming  chlorides  of  the  metals  and  liberat- 
ing hydrogen.    Unlike  nitric  and  sulfuric  acids,  it  has  no  oxidizing 
effects,  so  that  when  it  acts  upon  metals,  hydrogen  is  always  evolved. 

The  acid  also  acts  upon  oxides  and  hydroxides  of  the  metals,  con- 
verting them  into  the  corresponding  chlorides. 

2.  Action  with  oxidizing  agents.    Many  oxidizing  agents  act  upon 
hydrochloric  acid,  as  expressed  in  the  following  equation  : 


We  have  already  noted  that  advantage  is  taken  of  this  reaction  in  the 
preparation  of  chlorine. 

3.  Action  with  nitric  acid;  aqua  regia.  When  nitric  acid  acts  as  an 
oxidizing  agent,  it  usually  decomposes,  as  represented  in  the  following 
equation:  2HNO.=  H,O  +  2  NO  +  8O  (1) 

If  hydrochloric  acid  is  present,  the  oxygen,  as  fast  as  formed,  reacts 
with  the  acid  according  to  the  following  equation  : 

6  HC1  +  3  O  =  3  HO  +  3  Cl  (2) 


258  GENERAL   CHEMISTRY 

The  nitric  oxide  formed  according  to  equation  (1)  is  not  evolved  as 
such,  but  combines  with  the  chlorine  liberated  according  to  equation 
(2)  to  form  an  orange-yellow,  gaseous  compound  known  as  nitrosyl 
chloride  (NOC1)  :  2  NO  +  C12  =  2  NOC1  (3) 

By  combining  these  three  equations  in  the  regular  way  and  dividing 
the  resulting  equation  by  2,  in  order  to  get  its  simplest  form,  one 
obtains  the  following: 

HNO,  +  3  HC1  =  2  HO  +  NOC1  +  C10 

o  2.  2 

When  concentrated  nitric  and  hydrochloric  acids  are  mixed,  the  re- 
action expressed  in  the  above  equation  takes  place  slowly.  In  the 
presence  of  some  substance,  such  as  a  metal,  which  will  unite  with  the 
chlorine  as  fast  as  it  is  formed,  the  speed  of  the  reaction  is  greatly  in- 
creased. This  mixture  of  nitric  and  hydrochloric  acids  is  termed  aqua 
regia,  and  is  prepared  by  adding  1  volume  of  nitric  acid  to  3  volumes 
of  hydrochloric  acid.  It  acts  upon  metals  and  other  substances  more 
energetically  than  either  of  the  acids  separately,  and  owes  its  solvent 
power  not  to  its  acid  properties  but  to  the  action  of  the  chlorine 
which  is  liberated.  Consequently,  when  it  acts  upon  metals,  it  con- 
verts them  into  chlorides. 

It  may  be  added  that  this  mixture  was  well  known  to  the  alche- 
mists, who  termed  it  aqua  regia  because  of  its  strong  solvent  powers. 
It  is  evident  that  any  other  mixture  of  substances,  the  constituents  of 
which  interact  to  liberate  chlorine,  has  a  similar  solvent  power. 

Salts  of  hydrochloric  acid ;  chlorides.  The  chlorides  of  all  the  metals 
are  known,  and  many  of  them  are  very  important  compounds.  Some 
of  them,  as  sodium  chloride  and  potassium  chloride,  are  found  in 
nature.  A  number  of  the  metals,  including  copper,  mercury,  and 
tin,  combine  with  different  percentages  of  chlorine,  and  thus  form 
two  chlorides.  Nearly  all  the  chlorides  of  the  metals  are  solids  and, 
with  the  exception  of  silver,  mercurous,  lead,  and  thallous  chlorides, 
are  all  soluble  in  water.  The  insoluble  chlorides  may  be  formed  by 
the  general  method  employed  for  preparing  insoluble  compounds 
(p.  200).  Some  of  the  more  important  chlorides  are  the  following: 
sodium  chloride  (salt)  (NaCl) ;  potassium  chloride  (KC1) ;  mercurous 
chloride  (calomel)  (HgCl)  ;  mercuric  chloride  (corrosive  sublimate) 
(HgCl2)  ;  ferric  chloride  (FeClg)  ;  barium  chloride  (BaCl2) ;  calcium 
chloride  (CaCl). 


THE  CHLORIDE  FAMILY 


259 


BROMINE 

History  and  occurrence.  Bromine  occurs  in  nature  combined  with 
certain  metals,  principally  sodium,  potassium,  calcium,  and  magnesium. 
These  compounds  are  known  as  bromides.  Large  quantities  of  bro- 
mides are  found  in  the  famous  potash  deposits  at  Stassfurt,  Germany 
(p.  405).  They  also  occur  in  the  waters  of  many  springs  and  deep 
wells,  mixed  with  relatively  large  quantites  of  sodium  chloride.  When 
such  waters  are  evaporated,  the  sodium  chloride  separates  first,  since 
it  is  present  in  much  larger  quantities  and  is  less  soluble  than  the 
bromides.  The  liquor  remaining  after  the  separation  of  most  of  the 
sodium  chloride  is  known  as  the  mother  liquor,  and  contains  the  bro- 
mides in  solution.  It  w;as  from  this  liquor  that  the  German  chemist 
Liebig  first  isolated  bromine.  He  concluded,  however,  that  the  red 
liquid  which  he  obtained  was  simply  a  compound  of  chlorine  and 
iodine.  A  few  months  later  (1826)  the  French  chemist  Ballard  again 
obtained  the  substance  from  similar  liquors.  He  rightly  considered 
it  to  be  an  elementary  substance,  and,  because  of  its  disagreeable 
odor,  named  it  bromine,  a  word  meaning  "  stench." 

Preparation.  The  general  methods  used  in  the  preparation  of  chlo- 
rine may  likewise  be  employed  in  separating  bromine  from  its 
compounds.  The  laboratory  and 
commercial  methods  most  largely 
used  are  the  following : 

1.  Laboratory  method.  In  the 
laboratory  bromine  is  most  often 
prepared  by  the  action  of  oxidiz- 
ing agents  upon  hydrogen  bro- 
mide or  hydrobromic  acid.  Since 
hydrogen  bromide  is  unstable,  it 
is  more  convenient  to  generate 
it  in  the  course  of  the  reaction 
by  using  a  mixture  of  sodium  bromide  and  sulfuric  acid.  The  oxidiz- 
ing agent  generally  used  is  manganese  dioxide.  The  reaction  is  en- 
tirely similar  to  that  used  in  the  preparation  of  chlorine  (p.  248),  and 
is  expressed  in  the  following  equation : 

2  NaBr  +  3  H2SO4  +  MnO2  =  2  NaHSO4  +  MnSO4  +  2  H2O  +  Br2 

The  bromide  and  manganese  dioxide  are  thoroughly  mixed  and  are  then 
introduced  into  the  retort  A  (Fig.  95),  the  end  of  which  just  touches  the  water  in 


FIG.  95 


260  GENERAL  CHEMISTRY 

the  flask  B.  The  sulfuric  acid  is  then  added.  As  the  retort  is  gently  heated,  the 
bromine  is  liberated,  distills  over,  and  collects  under  the  water  in  the  flask.  The 
latter  is  kept  cool  by  immersion  in  ice  water  in  the  beaker  C. 

In  addition  to  the  general  methods  applicable  alike  to  the  prepara- 
tion of  chlorine  and  bromine,  an  additional  method  may  be  used  for 
the  preparation  of  bromine.  This  method  is  based  upon  the  fact  that 
chlorine  readily  liberates  bromine  from  its  compounds  with  the  metals : 

2  NaBr  +  Cla  =  2  NaCl  +  Br2 

This  action  is  quite  similar  to  the  displacement  of  one  metal  by  another  which 
precedes  it  in  the  electromotive  series  ;  thus, 

CuSO4  +  Zn  =  ZnSO4  +  Cu 

The  nonmetals,  as  well  as  the  metals,  can  be  arranged  in  an  electromotive 
series,  and  in  such  an  arrangement  the  halogen  elements  occur  in  the  order 
F  —  Cl  —  Br  —  I.  Any  one  of  these  is  displaced  from  its  salts  in  solution  by  all 
those  which  precede  it. 

2.  Commercial  method.  In  the  United  States,  bromine  is  obtained 
commercially  from  salt  water,  and  Michigan  furnishes  by  far  the 
largest  quantity.  Smaller  amounts  are  obtained  from  salt  waters 
taken  from  deep  wells  along  the  Ohio  River.  In  Michigan  the  bromine 
is  separated  from  the  salt  water  by  electrolysis.  Some  chlorine  is  also 
set  free  along  with  the  bromine,  but  this  reacts  with  the  bromides 
present  in  solution,  forming  chlorides  and  liberating  bromine,  as  ex- 
plained in  the  previous  paragraph.  In  the  Ohio  River  valley,  on  the 
other  hand,  the  bromine  is  obtained  by  treating  the  mother  liquors 
resulting  from  the  removal  of  the  salt  with  sulfuric  acid  and  sodium 
chlorate.  The  acid  reacts  with  the  bromides,  forming  hydrobromic 
acid,  which  is  oxidized  to  free  bromine  by  the  oxygen  resulting  from 
the  sodium  chlorate.  The  equations  are  as  follows  : 

NaBr  +  H0SO4  =  NaHSO4  +  HBr 
4  HBr  +  O2  =  2  H2O  +  2  Bra 

The  electrolytic  method  has  the  advantage  of  not  necessitating  the 
removal  of  the  salt  from  the  waters  in  order  to  obtain  the  bromine. 

•  In  Europe  the  source  of  bromine  is  the  mother  liquors  left  in  the 
process  of  separating  certain  salts  occurring  in  the  Stassfurt  deposits. 
It  is  obtained  from  these  by  the  general  methods  described  above. 

Instead  of  shipping  the  bromine  in  the  liquid  state  it  is  sometimes 
found  convenient  to  form  a  bromide  of  iron,  which  is  more  easily 
transported.  Bromine  is  readily  liberated  as  desired. 


THE  CHLOKINE  FAMILY  261 

Properties.  Bromine  is  a  dark-red  liquid  whose  density  is  3.102.  Its 
vapor  has  an  offensive  odor  and  is  very  irritating  to  the  eyes  and 
throat.  The  liquid  boils  at  63°  and  solidifies  at  —7°,  but  even  at  ordi- 
nary temperatures  it  has  a  high  vapor  pressure,  so  that  it  evaporates 
rapidly,  forming  a  reddish-brown  gas  very  similar  to  nitrogen  dioxide  in 
appearance.  At  20°,  100  volumes  of  water  dissolves  about  1  volume 
of  bromine,  forming  a  reddish  Solution  called  bromine  water.  Bromine 
is  readily  soluble  in  carbon  disulfide,  forming  a  reddish  solution. 

Chemical  conduct.  The  chemical  conduct  of  bromine  is  very  similar 
to  that  of  chlorine,  except  that  it  is  less  active.  It  combines  directly 
with  many  of  the  same  elements  with  which  chlorine  unites,  but  with 
less  energy.  It  combines  with  hydrogen,  and  even  abstracts  it  from 
some  of  its  compounds.  As  would  be  expected,  its  bleaching  action  is 
much  less  marked  than  that  of  chlorine.  Its  solution  in  water  is  often 
used  as  an  oxidizing  agent.  For  example,  sulfurous  acid  is  readily 
converted  into  sulfuric  acid  by  the  addition  of  a  suitable  amount  of 
bromine  water:  Bri+  HfO  =  2  HBr  +  O 


Uses.  Bromine  is  used  principally  in  the  preparation  of  bromides, 
which  are  employed  to  a  considerable  extent  in  photography  and  as 
medicinal  agents.  It  is  likewise  used  in  the  preparation  of  a  number 
of  organic  drugs  and  dyestuffs. 

Hydrogen  bromide.  One  would  naturally  expect  that  hydrogen  bro- 
mide could  be  prepared  by  the  same  general  method  as  that  employed 
in  the  preparation  of  hydrogen  fluoride  and  hydrogen  chloride,  namely, 
by  the  action  of  sulfuric  acid  upon  a  bromide  such  as  NaBr  or  KBr  : 

NaBr  +  H2SO4  =  NaHSO4  +  HBr 

This  reaction  does  indeed  take  place,  the  hydrogen  bromide  being 
evolved  in  the  form  of  a  colorless  gas  which  fumes  strongly  in  the 
air.  At  the  same  time  some  bromine  is  liberated,  as  is  indicated  by  the 
formation  of  a  reddish  vapor.  The  odor  of  sulfur  dioxide  can  also  be 
detected.  This  difference  in  the  action  of  sulfuric-  acid  upon  fluorides 
and  chlorides,  on  the  one  hand,  and  upon  bromides,  on  the  other,  is  due 
to  the  relatively  unstable  character  and  consequent  reducing  proper- 
ties of  hydrogen  bromide.  In  the  presence  of  concentrated  sulfuric 
acid,  which  is  a  good  oxidizing  agent,  a  portion  of  the  hydrogen  bro- 
mide formed  is  decomposed,  the  bromine  being  liberated  while  the 


262 


GENERAL  CHEMISTRY 


hydrogen  is  oxidized  to  water.    The  sulfuric  acid  is  reduced  to  sul- 
furous  acid  in  the  process : 


2  HBr  +  H(>S04  =  H2SO 


H2Q 


Br 


This  method,  therefore,  cannot  be  used  for  the  preparation  of  pure 
hydrogen  bromide. 

The  method  usually  employed  in  the  preparation  of  hydrogen 
bromide  consists  in  the  action  of  water  upon  phosphorus  tribromide. 
The  latter  compound  is  a  colorless  liquid  formed  by  the  union  of 
phosphorus  and  bromine,  and  has  the  formula  PBr3.  When  brought  in 
contact  with  water,  it  undergoes  complete  hydrolysis  (p.  226),  form- 
ing hydrogen  bromide  and  phosphorous  acid  (H3PO3  or  P(OH)3). 
This  reaction  is  made  clearer  by  the  use  of  structural  formulas  : 


OH 


OH 


OH  =  3  HBr  +  P~OH 
OH  XOH 


The  preparation  is  carried  out  as  follows :  Some  red  phosphorus  is  introduced 
into  a  flask  A  (Fig.  96),  and  sufficient  water  is  added  to  cover  it.  The  separatory 
funnel  B  contains  the  bromine.  By  means  of  the  stopcock,  bromine  is  allowed 

to  flow  drop  by  drop  from  the  funnel 
into  the  flask.  The  bromine,  on  coming 
in  contact  with  the  phosphorus,  com- 
bines with  it  to  form  phosphorus  tri- 
bromide, which  then  reacts  with  water. 
The  equations  are 


2  P  +  3  Br,  =  2  PBr, 


(1) 


PBr3  +  3  H20  -  P  (OH),  +  3  HBr    (2) 

The  U-tube  C  contains  glass  beads 
which  have  been  moistened  with  water 
and   rubbed  in  red  phosphorus.    Any 
FIG.  96  bromine  escaping   action  in   the  flask 

acts  upon  the  phosphorus  in  the  U-tube. 

The  hydrogen  bromide  is  collected  in  D  by  displacement  of  air.  An  aqueous 
solution  of  the  gas  can  be  prepared  in  the  same  way  as  an  aqueous  solution  of 
hydrogen  chloride. 

Properties.  Hydrogen  bromide  very  strikingly  resembles  hydrogen 
chloride  in  its  properties.  It  is  a  colorless,  strongly  fuming  gas  and 
may  be  condensed  to  a  colorless  liquid  which  boils  at  -  69°.  It  is 
very  soluble  in  water.  Under  standard  conditions  1  volume  of 
water  dissolves  612  volumes  of  the  gas.  The  resulting  solution  has 


THE  CHLORINE  FAMILY  263 

a  density  of  1.5  and  contains  88  per  cent  of  the  gas.  An  aqueous 
solution  containing  48  per  cent  of  hydrogen  bromide  boils  at  126° 
and  distills  with  unchanged  concentration. 

Chemical  conduct  of  hydrobromic  acid.  Hydrogen  bromide,  like  hy- 
drogen chloride,  has  but  little  activity.  When  dissolved  in  water,  it 
dissociates  into  the  ions  H+  and  Br~,  so  that  the  solution  is  strongly 
acid  and  is  known  as  hydrobromic  acid.  It  is  very  similar  to  hydro- 
chloric acid.  It  reacts  with  metals,  and  with  their  oxides  and  hydrox- 
ides, forming  the  corresponding  bromides.  It  differs  from  hydrochloric 
acid  mainly  in  that  it  is  much  more  easily  oxidized,  so  that  bromine  is 
more  readily  liberated  from  it  than  chlorine  is  from  hydrochloric  acid. 
It  is  therefore  a  moderately  active  reducing  agent.  Free  chlorine  acts 
upon  hydrobromic  acid,  liberating  bromine,  as  represented  in  the  fol- 
lowing  equation:  2  HBr  +  C12  =  2  HC1  +  Br, 

Salts  of  hydrobromic  acid ;  bromides.  The  bromides  are  in  general 
very  similar  to  the  chlorides  in  their  properties,  and  are  prepared  by 
the  same  general  methods.  They  are  all  soluble  except  silver  bromide, 
mercurous  bromide,  and  lead  bromide  (compare  chlorides,  p.  258). 
Silver  bromide  is  used  in  photography,  while  sodium  bromide  and 
potassium  bromide  are  used  as  medicinal  agents. 

IODINE 

History  and  occurrence.  Iodine  is  present  in  sea  water,  but  in  rela- 
tively small  quantities.  Certain  seaweeds  absorb  the  iodine  from  the 
water,  thus  concentrating  it  within  then-  tissues.  It  was  from  the  ashes 
obtained  by  burning  seaweed  that  the  French  chemist  Courtois,  in 
1812,  first  isolated  the  element,  which  he  termed  iodine  (from  the 
Greek  word  meaning  "  violet-colored ")  because  of  the  violet  color 
of  its  vapor.  Iodine  is  also  found  in  certain  animal  life  of  the  sea, 
such  as  sponges,  oysters,  and  some  fishes.  It  likewise  occurs  in  the 
deposits  of  Chile  saltpeter  (sodium  nitrate),  and  this  at  present 
constitutes  the  largest  source  of  commercial  iodine.  It  is  interesting 
to  note  that  small  amounts  of  iodine  exist  in  the  human  body  in 
the  thyroid  gland. 

Preparation.  The  principal  methods  used  in  the  preparation  of  iodine 
are  the  following : 

1.  Laboratory  method.  Iodine  is  liberated  from  the  iodides  by  the 
action  of  sulfuric  acid  and  manganese  dioxide.  The  reaction  is  similar 


GENERAL  CHEMISTRY 

to  that  which  takes  place  in  the  liberation  of  chlorine  from  the  chlorides, 
and  of  bromine  from  the  bromides.    The  equation  is  as  follows : 

2  Nal  +  3  H2S04  +  MnO2  =  2  NaHSO4  +  MnSO4  +  2  H2O  + 12 

This  method  serves  as  a  convenient  one  for  the  preparation  of  iodine  in  the 
laboratory.  The  apparatus  is  the  same  as  that  used  in  the  preparation  of  bromine 
(Fig.  95,  p.  259).  A  mixture  of  manganese  dioxide  and  sodium  or  potassium 
iodide  is  placed  in  the  retort  A,  sulfuric  acid  added,  and  a  gentle  heat  applied. 
The  iodine  is  evolved  in  the  form  of  a  violet-colored  vapor  which  condenses  to  a 
purplish-black  crystalline  solid  on  the  colder  portions  of  the  retort.  By  regula- 
ting the  heat  it  can  be  driven  over  and  condensed  in  the  flask  B,  which  is  kept 
cool  by  ice  water. 

2.  Commercial  method.  Commercial  iodine  is  obtained  either  from 
seaweeds  or  from  crude  Chile  saltpeter  (NaNO3),  which  is  known  as 
caliche. 

(a)  Preparation  from  caliche.    The  iodine  is  distributed  through  the 
caliche  in  the  form  of  sodium  iodate  (NaIO3)  and  is  obtained  from 
the  mother  liquors  left  in  the  purification  of  the  nitrate.   The  iodine  is 
liberated  from  the  sodium  iodate  by  the  action  of  the  sulfites  of  sodium : 

2  NaI08  +  3  Na2S08  +  2  NaHSO8'=  5  Na2SO4  + 12  +'  H2O 

The  resulting  iodine  is  removed  by  filtration,  dried,  and  purified  by 
sublimation. 

(b)  Preparation  from  seaweeds.    Previous  to  the  discovery  of  iodine 
in  the  Chile  saltpeter,  the  element  was  obtained  entirely  from  sea- 
weeds.   These  weeds,  known  as  kelp,  were  collected  upon  the  shores 
of  Scotland,  Ireland,  Japan,  and  France,  and  were  dried  and  burned. 
The  ashes,  also  known  as  kelp,  contain  a  number  of  compounds  of 
sodium,  especially  sodium  carbonate  and  chloride,  together  with  about 
0.3  per  cent  of  sodium  iodide.    The  mother  liquor  left  after  the  re- 
moval of  the  carbonate  and  chloride   contains   the  sodium  iodide. 
From  this  the  iodine  was  obtained  either  by  the  action  of  manganese 
dioxide  and  sulfuric  acid,  as  explained  above,  or  by  the  action  of 
chlorine:  2  Nal  +  C12  =  2  NaCl  + 12 

After  the  discovery  of  iodine  in  Chile  saltpeter,  the  production  of  the  element 
from  seaweeds  practically  ceased  for  a  time,  since  it  could  be  obtained  from  the 
saltpeter  at  a  much  lower  cost.  Later,  however,  the  method  of  recovering  the 
element  from  seaweeds  has  been  improved,  so  that  a  limited  amount  is  again 
obtained  from  this  source.  The  supply  of  iodine,  however,  is  greater  than  the 
demand,  and  new  uses  for  the  element  are  being  sought. 


THE  CHLORINE  FAMILY  265 

Purification  of  iodine.  Iodine  can  be  purified  very  conveniently  in  the  following 
way  :  The  crude  iodine,  mixed  intimately  with  a  little  potassium  iodide,  is  placed 
in  a  beaker  A  (Fig.  97),  in  the  top  of  which  rests  a  round-bottomed  flask  B, 
containing  cold  water.  The  apparatus  is  placed  upon  ^_^ 

a  sand  bath  C  and  gently  heated.  The  iodine  rapidly 
evaporates  and  condenses  again  on  the  cold  surface  of  /  \ 

the  flask  in  shining  crystals.  The  crude  iodine  often 
contains  small  amounts  of  free  chlorine  and  bromine. 
These  react  with  the  potassium  iodide  present,  forming 
respectively  potassium  chloride  and  potassium  bromide, 
with  the  corresponding  evolution  of  iodine.  In  this  way 
the  iodine  is  separated  from  all  nonvolatile  matter. 


Properties.  Iodine  is  a  purplish-black  shin- 
ing solid  which,  when  sublimed,  crystallizes  in 
brilliant  plates  belonging  to  the  rhombic  sys- 
tem. It  has  a  density  of  4.95,  melts  at  113°, 
and  boils  at  184.4°.  The  element  has  a  strong, 

unpleasant  odor,  although  not  so  disagreeable  as  that  of  chlorine  or 
bromine.  Even  at  ordinary  temperatures  it  gives  off  a  beautiful 
violet  vapor,  which,  increases  in  amount  as  heat  is  applied.  It  is  only 
slightly  soluble  in  water,  1  part  being  soluble  in  3750  parts  of  water 
at  15°.  It  is  more  readily  soluble  in  a  solution  of  potassium  iodide  or 
of  hydrogen  iodide,  forming  a  dark-brown  liquid.  It  also  dissolves 
in  carbon  disulfide,  forming  a  violet-colored  liquid.  Its  solution  in 
alcohol  is  known  as  tincture  of  iodine  and  is  used  in  medicine.  When 
applied  to  the  skin,  it  produces  a  brown  stain. 

Chemical  conduct.  Iodine  is  similar  to  chlorine  and  bromine  in  its 
chemical  properties,  but  is  less  active.  Both  of  the  latter  elements 
liberate  iodine  from  its  compounds  with  hydrogen  or  the  metals: 

+  C12=2HC1  +  I2 
+  Cl=2NaCl  +  I 


Like  chlorine  and  bromine,  it  combines  directly  with  many  of  the 
metals  as  well  as  with  the  nonmetals.  In  the  presence  of  water  it 
acts  as  a  mild  oxidizing  agent. 

A  very  characteristic  property  of  iodine  is  its  power  of  imparting 
a  blue  color  to  a  solution  of  starch.  The  reaction  is  a  very  delicate 
one,  as  can  be  shown  by  adding  a  few  drops  of  an  aqueous  solu- 
tion of  iodine  to  a  test  tube  containing  starch  solution.  The  blue 
color  of  the  resulting  solution  fades  on  heating,  but  forms  again  as 
the  solution  cools. 


266  GENERAL  CHEMISTRY 

The  cause  of  the  production  of  this  blue  color  is  not  known  with  certainty. 
By  some  it  is  regarded  as  due  to  the  formation  of  an  unstable  compound  of  iodine 
and  starch,  which  is  blue  in  color.  When  heated,  this  compound  dissociates  into 
its  constituents  ;  hence  the  color  fades.  On  cooling,  they  recombine  and  the  color 
again  appears.  Others  regard  the  color  as  due  simply  to  the  formation  of  a 
solution  of  the  iodine  in  the  starch. 

The  production  of  the  blue  color  serves  as  a  very  delicate  test 
either  for  free  iodine  or  for  starch.  The  color  is  not  produced  by 
compounds  of  iodine.  This  can  be  shown  by  adding  a  few  drops  of 
a  solution  of  potassium  iodide  to  a  starch  solution.  No  apparent 
change  takes  place.  If  some  chlorine  water  is  now  added  to  the 
mixture,  the  iodine  is  liberated  and  the  blue  color  at  once  appears. 

Uses  of  iodine.  Iodine  is  used  extensively  in  medicine,  especially 
in  the  form  of  tincture  of  iodine.  It  is  also  used  in  the  preparation  of 
the  iodides  and  of  certain  organic  dyes  and  drugs.  The  common 
antiseptic  known  as  iodoform  has  the  formula  CHI3. 

Hydrogen  iodide.  The  method  generally  employed  for  the  prepara- 
tion of  hydrogen  iodide  is  similar  to  that  used  for  the  preparation 
of  hydrogen  bromide  (p.  262)  and  consists  in  the  reaction  between 
phosphorus  tri-iodide  and  water: 

OH  .OH 

3HI  +  P-OH 
OH  XOH 

The  hydrogen  iodide  is  evolved  as  a  heavy  colorless  gas  and  may 
be  collected  by  the  displacement  of  air.  An  aqueous  solution  of  the 
gas  can  be  prepared  by  passing  hydrogen  sulfide  into  water  containing 
finely  divided  iodine  in  suspension  : 


When  the  reaction  is  complete,  the  precipitated  sulfur  is  removed  by 
filtration.  Solutions  of  the  gas  can  be  prepared  in  this  way  up  to 
50  per  cent  strength. 

It  will  be  recalled  that  hydrogen  bromide,  because  of  its  unstable  character  and 
the  consequent  ease  with  which  it  is  oxidized,  cannot  be  prepared  in  the  pure  state 
by  the  action  of  sulfuric  acid  on  the  bromides.  Since  hydrogen  iodide  is  more 
unstable  than  hydrogen  bromide,  it  is  evident  that  this  general  method  is  still 
less  adapted  to  its  preparation. 

The  reactions  which  take  place  when  sulfuric  acid  is  added  to  an  iodide  are 
expressed  in  the  following  equations  : 

Nal  +  H2SO4  =  NaHSO4  +  HI  (1) 

8  HI  +  H2S04  =  4  H20  +  H2S  +  4  12  (2) 


THE  CHLORINE  FAMILY  267 

The  hydrogen  sulfide  in  equation  (2)  may  react  with  any  excess  of  sulfuric  acid 
to  form  sulfurous  acid  and  free  sulfur  (p.  209). 

It  will  be  noted  from  the  above  equations  that  the  reduction  of  the  sulfuric 
acid  is  more  complete  than  in  the  similar  reactions  between  sulfuric  acid  and 
sodium  bromide  (p.  262).  In  the  latter  case  the  acid  is  reduced  simply  to 
H2SO3,  while  in  its  action  on  an  iodide  all  of  its  oxygen  is  given  up,  the  acid 
being  reduced  thereby  to  H2S. 

Properties  of  hydrogen  iodide.  Hydrogen  iodide  resembles  hydrogen 
chloride  and  hydrogen  bromide  in  its  physical  properties,  being  a 
strongly  fuming  colorless  gas.  It  is  4.37  times  as  heavy  as  air.  At 
0°  it  is  condensed  to  a  colorless  liquid  by  a  pressure  of  4  atmospheres. 
At  10°  about  450  volumes  of  the  gas  dissolves  in  1  volume  of  water. 
A  solution.  containing  57  per-  cent  of  hydrogen  iodide  boils  at  127° 
and  distills  with  unchanged  concentration.  Owing  to  the  ease  with 
which  the  gas  is  decomposed  into  its  elements,  it  acts  in  many 
respects  like  hydrogen,  being  a  strong  reducing  agent.  This  might 
be  expected  from  the  fact  that  it  is  an  endothermic  compound,  as  shown 
in  the  equation  = 


In  an  atmosphere  of  oxygen  it  burns,  forming  water  and  iodine. 
Pure  hydrogen  iodide,  whether  in  the  form  of  a  gas  or  of  a  liquid, 
is  neutral.  When  dissolved  in  water,  it  ionizes  as  follows: 

HI  ^=>  H+  +  1- 

This  solution  has  strong  acid  properties  due  to  the  hydrogeu  ions 
present,  and  is  known  as  hydriodic  acid. 

Chemical  conduct  of  hydriodic  acid.  Hydriodic  acid  differs  from 
hydrochloric  and  hydrobromic  acid  mainly  in  the  ease  with  which  it 
is  oxidized.  The  freshly  prepared  solution  is  colorless,  but  soon  turns 
brown,  owing  to  the  liberation  of  iodine  by  the  oxygen  of  the  air: 


As  the  action  continues,  the  iodine  separates  in  crystalline  form.  The 
acid,  as  well  as  hydrogen  iodide,  is  therefore  a  strong  reducing  agent. 

Hydriodic  acid  reacts  with  many  of  the  metals,  as  well  as  with  their 
oxides  and  hydroxides,  forming  the  corresponding  salts. 

Salts  of  hydriodic  acid  ;  iodides.  These  compounds  are  similar  to  the 
corresponding  chlorides  and  bromides,  but  are  not  so  stable  toward 
heat.  They  are  all  solids  and,  with  the  exception  of  the  iodides 
of  silver,  mercury,  and  lead,  are  soluble  in  water.  Silver  iodide  is  used 
in  photography  and  potassium  iodide  in  medicine. 


CHAPTER  XX 

THE  OXYGEN  COMPOUNDS  OF  THE  HALOGENS 

General.  While  neither  chlorine  nor  iodine  combines  with  oxygen 
directly,  nevertheless  a  number  of  oxides  of  these  two  elements  have 
been  prepared  by  indirect  methods,  as  described  below.  Fluorine  and 
bromine,  on  the  other  hand,  do  not  form  oxides.  With  the  exception 
of  fluorine,  the  halogens  form  oxygen-  acids.  The  salts  •  of  some  of 
these  acids  are  of  considerable  importance. 

The  oxides  and  oxygen  acids  of  chlorine.  The  following  table  includes 
the  names  and  formulas  of  the  oxides  and  oxygen  acids  of  chlorine, 
and  also  shows  their  relation  to  each  other. 

OXIDES  ACIDS 

chlorine  monoxide  rH2O  +  C12O    =  2  HC1O, 


Cl  O 

2      1  (anhydride  of  hypochlorous  acid)  /  \      hypochlorous  acid 

f  chlorine  trioxide  (unknown)  j  f  H2O  +  C12O3  =  2  HC1O2 

2    8  \  (anhydride  of  chlorous  acid)  /  \     chlorous  acid 

0  |  chlorine  pentoxide  (unknown)  ^  J  H2O  +  C12O5  =  2  HC1O3, 


2    6\  (anhydride  of  chloric  acid)        J  \     chloric  acid 

C1  o   f  chlorine  heptoxide  j  f  H2O  +  C12O7  =  2  HC1O4, 

2    7  \  (anhydride  of  perchloric  acid)  /  \     perchloric  acid 
C1O2      chlorine  dioxide  (peroxide). 

The  oxides  of  chlorine.  The  three  known  oxides  of  chlorine,  namely, 
C12O,  C12O7,  and  C1O2,  are  all  unstable  compounds.  A  brief  discussion 
of  each  follows  : 

1.  Chlorine  monoxide  (hypochlorous  anhydride)  (C720).  This  compound 
is  prepared  by  passing  chlorine  through  a  tube  containing  mercuric 

2  C12  +  2  HgO  =  HgO  •  HgCl2  +  C12O 

It  is  a  highly  explosive  yellow  gas,  which  condenses  to  a  liquid  at 
5°.    With  water  it  forms  hypochlorous  acid.^ 

2.  Chlorine  heptoxide  (perchloric  anhydride)   (C/207).   This   oxide   is 
formed  by  the  action  of  a  strong  dehydrating  agent,  such  as  phos- 
phorus pentoxide  (P2O6),  on  perchloric  acid: 


268 


THE  OXYGEN  COMPOUNDS  OF  THE  HALOGENS      269 

It  is  a  colorless  oily  liquid  and  explodes  with  great  violence  when 
ignited  or  struck.  With  water  it  forms  perchloric  acid. 

3.  Chlorine  dioxide  (C/02).  This  oxide  results  from  the  decomposition 
of  chloric  acid,  as  is  represented  in  the  following  equation : 

3  HC1O3  =  HC1O4  +  H2O  +  2  C1O2 

It  is  prepared  by  the  action  of  sulfuric  acid  upon  potassium  chlorate. 
Chloric  acid  is  first  formed,  but  immediately  decomposes  according 
to  the  above  equation.  The  reaction  must  be  carried  out  with  great 
care ;  otherwise  the  decomposition  may  take  place  with  explosive  vio- 
lence. The  intensity  of  the  reaction  may  be  shown  by  touching  a 
small  crystal  of  potassium  chlorate  with  a  glass  rod  moistened  with 
concentrated  sulfuric  acid. 

Chlorine  dioxide  is  a  yellow  gas  which  may  be  condensed  to  a  liquid 
boiling  at  10°.  Just  as  nitrogen  dioxide  (NO2)  reacts  with  water  to 
form  a  mixture  of  nitrous  and  nitric  acids,  so  chlorine  dioxide,  under 
the  same  conditions,  forms  a  mixture  of  chlorous  and  chloric  acids. 
The  similarity  between  the  two  reactions  is  shown  in  the  following 
equations :  2  NO2  +  H2O  =  HNO2  +  HNO3 

2  C102  +  H20  =  HC102  +  HC103 

Hypochlorous  acid  and  the  hypochlorites.  Both  the  free  hypochlorous 
acid  and  its  salts,  namely,  the  hypochlorites,  are  unstable  and  have 
only  been  obtained  in  dilute  solution. 

1.  Preparation.  Hypochlorous  acid  can  most  readily  be  obtained 
from  its  salts.  Solutions  of  the  hypochlorites  of  sodium,  potassium, 
and  calcium  are  formed,  along  with  their  chlorides,  by  passing  chlorine 
into  cold  solutions  of  their  respective  hydroxides.  Thus,  with  potassium 
hydroxide  the  reaction  is  expressed  by  the  following  equation :  . 

2  KOH  +  C12  =  KC1O  +  KC1  +  H2O 

From  the  resulting  hypochlorites  a  solution  of  hypochlorous  acid  can 
be  prepared  by  adding  just  sufficient  dilute  sulfuric  acid  to  react  with 
the  hypochlorite,  as  expressed  in  the  following  equation : 

2  KC1O  -f  H2SO4  =  K2SO4  +  2  HC1O 

On  distilling  the  resulting  mixture,  a  solution  of  hypochlorous  acid  in 
water  is  obtained.  Dilute  solutions  of  hypochlorous  acid  can  also  be 
obtained  by  the  action  of  chlorine  monoxide  upon  water. 


270  GENERAL  CHEMISTRY 

2.  Properties.  Both  hypochlorous  acid  and  the  hypochlorites  are 
excellent  oxidizing  agents.  In  the  presence  of  a  substance  that  will 
combine  with  the  oxygen  formed,  they  decompose  as  follows  : 

HC1O  =  HC1  +  O  KC1O  =  KC1  +  O 

On  the  other  hand,  when  their  solutions  are  heated,  hypochlorous 
acid  and  its  salts  form  chloric  acid  and  chlorates  respectively: 


IKCI  o 

=  HC1O8  +  2  HCl  KCllO 

KC1  O 


=  KC1O,  +  2  KC1 


It  is  evident,  therefore,  that  if  one  wishes  to  prepare  hypochlorites,  the 
solutions  must  be  kept  cold  ;  otherwise  chlorates  are  obtained. 

Uses.  The  hypochlorites,  as  well  as  the  free  acid,  are  used  as  oxidiz- 
ing agents,  especially  in  bleaching  (p.  253). 

Bleaching  powder  (CaOCl2).  When  chlorine  is  passed  into  a  cold 
solution  of  calcium  hydroxide,  there  is  formed  a  mixture  of  the 
chloride  and  hypochlorite  of  calcium  (p.  269)  ;  if  passed  over  the  dry 
calcium  hydroxide,  however,  there  is  formed  a  white  solid  compound 
known  commercially  as  bleaching  powder  or  chloride  of  lime  : 

Ca(OH)2  +  C12  =  CaOCl2  +  H2O 

The  reactions  of  this  compound  are  best  explained  on  the  assumption 
that  it  has  the  structural  formula  Ca<^|-;1.  In  accordance  with  this 
formula  it  must  be  regarded  as  a  mixed  salt,  namely,  a  calcium  salt 
of  hypochlorous  and  hydrochloric  acids,  being  formed  by  the  dis- 
placement of  one  atom  of  hydrogen  in  a  molecule  of  each  of  these 
acids  by  a  divalent  calcium  atom.  When  an  acid  such  as  sulfuric  is 
added  to  bleaching  powder,  free  hypochlorous  and  hydrochloric  acids 
are  liberated  and  react  with  each  other  to  form  water  and  chlorine  : 

OPl 


HC1 

HC10  +  HC1  =  H20  +  C12 

When  bleaching  powder  is  exposed  to  air,  hypochlorous  acid  is  liber- 
ated through  the  action  of  moisture  and  carbon  dioxide. 

Uses  of  bleaching  powder.  Bleaching  powder  is  made  in  large  quantities  from 
chlorine  obtained  by  the  electrolysis  of  sodium  chloride  and  is  used  commer- 
cially as  a  source  of  chlorine,  since  it  is  easily  prepared  and  transported  and  the 
chlorine  present  can  be  liberated  as  desired.  The  commercial  product  generally 
contains  from  35  to  3  7  per  cent  of  available  chlorine. 


THE  OXYGEN  COMPOUNDS  OF  THE  HALOGENS      271 

Chlorous  acid  and  the  chlorites.  Chlorous  acid  is  formed  in  small 
quantities  when  chlorine  dioxide  is  dissolved  in  water: 

2  C102  +  H20  =  HC102  +  HC103 

If  this  solution  is  neutralized  with  potassium  hydroxide,  a  mixture  of 
potassium  chlorite  and  chlorate  results.  The  chlorites  of  a  few  of  the 
other  metals  are  known,  but  they  are  all  very  unstable. 

Chloric  acid  and  the  chlorates.  The  chlorates  of  the  metals  that  form 
soluble  hydroxides  are  prepared  by  passing  chlorine  into  hot  solutions 
of  their  respective  hydroxides,  as  already  explained  (p.  270).  With 
potassium  hydroxide  the  reactions  are  expressed  by  the  following 
equations :  Cl,  +  2  KOH  =  KC1O  +  KC1  +  H2O  ''  . 

3  KC1O  =  KC1O3  +  2  KC1 

By  combining  the  two  equations  the  following  is  obtained : 
3  C12  +  6  KOH  =  KC1O3  +  5  KC1  +  H2O 

When  the  resulting  solution  is  evaporated,  the  potassium  chlorate, 
being  much  less  soluble  than  the  potassium  chloride,  separates  first, 
and  by  repeated  crystallization  can  be  obtained  in  a  pure  state. 

From  the  chlorates,  chloric  acid  itself  can  be  prepared.  The  most 
convenient  method  consists  in  adding  sulfuric  acid  to  an  aqueous 
solution  of  barium  chlorate : 

Ba(C103)2  +  H2S04  =  BaS04  +  2  HC1O3 

The  barium  sulfate,  being  insoluble,  separates  as  a  white  precipitate, 
which  is  removed  by  filtration,  leaving  a  solution  of  chloric  acid.  This 
may  be  concentrated  until  it  contains  40  per  cent  of  acid.  Further 
concentration  leads  to  the  decomposition  of  the  acid,  forming  per- 
chloric acid,  water,  and  chlorine  dioxide.  The  concentrated  aqueous 
solution  of  the  acid  is  a  colorless  liquid  and  has  powerful  oxidizing 
properties. 

The  chlorates  can  readily  be  obtained  in  the  pure  state.  They  are 
all  soluble  in  water  and  ionize  as  follows : 

KC1O3^=±:K+,  C1O3- 

Potassium  chlorate  is  perhaps  the  most  important  of  these  salts,  being 
used  in  the  preparation  of  oxygen  and  as  an  oxidizing  agent. 

Preparation  of  hypochlorites  and  chlorates  by  electrolytic  methods.  It  will  be 
recalled  that  the  electrolysis  of  solutions  of  potassium  chloride  or  of  sodium 
chloride  results  in  the  formation  of  chlorine,  together  with  the  corresponding 


272 


GENERAL  CHEMISTRY 


hydroxides  of  the  metals.  It  is  possible  to  so  regulate  this  process  that  the 
chlorine,  instead  of  being  evolved,  is  retained  in  the  solution,  together  with  the 
hydroxides,  with  which  it  interacts  to  form  hypochlorites  or  chlorates,  according 
to  the  equations  given  above.  This  method  is  now  coming  into  general  use  for 
the  preparation  of  these  salts.  It  is  possible  to  obtain  either  the  hypochlorites  or 
chlorates  by  properly  choosing  the  conditions  of  the  electrolysis. 

Perchloric  acid  and  the  perchlorates.    When  potassium  chlorate  is 
heated,  a  portion  of  the  compound  changes  into  the  perchlorate  : 

4  KC1CX  =  3  KC1O4  +  KC1 

o  4 

This  reaction  serves  as  a  convenient  method  for  preparing  perchlorates, 
and  perchloric  acid  itself  can  be  obtained  from  these  by  the  addition  of 
sulfuricacid:  ' 


The  perchloric  acid  formed  is  separated  by  distillation.  This  process, 
however,  cannot  be  carried  on  under  atmospheric  pressure,  since  the 
temperature  required  decomposes  the  acid.  Under  greatly  diminished 
pressure  the  boiling  point  of  the  acid  is  lowered  to  such  an  extent 
that  the  compound  may  be  distilled  without  decomposition. 


ft 


FIG.  98 


To  Air  Pump 


To  distill  a  liquid  under  less  than  atmospheric  pressure,  the  apparatus  repre- 
sented in  Fig.  98  may  be  used.  The  liquid  is  placed  in  the  distilling  flask  A,  the 
delivery  tube  of  which  is  connected  with  the  condenser  B.  This  is  in  turn  con- 
nected by  a  rubber  stopper  with  the  strong  receiving  flask  C,  the  neck  of  which 
is  provided  with  a  side  tube  Z),  which  is  attached  to  an  air  pump.  A  manometer 
E,  for  indicating  the  pressure,  may  be  inserted  between  the  receiving  flask  and 
the  air  pump.  After  exhausting  the  air  in  the  apparatus  to  the  desired  extent, 
the  liquid  in  A  is  gently  heated  and  distilled  over  into  C. 


THE  OXYGEN  COMPOUNDS  OF  THE  HALOGENS      273 

Pure  perchloric  acid  is  a  colorless  liquid.  It  is  unstable,  sometimes 
decomposing  spontaneously  with  great  violence.  Like  the  other  oxygen 
acids  of  chlorine,  it  is  an  excellent  oxidizing  agent.  The  perchlorates 
can  be  obtained  in  the  pure  state  and  are  the  most  stable  of  all  the 
salts  of  the  oxygen  acids'  of  chlorine.  They  are  soluble  in  water, 
forming  the  ion  C1O4~,  together  with  the  metal  ion.  At  high  tempera- 
tures they  decompose  into  oxygen  and  the  corresponding  chlorides. 

The  oxygen  acids  of  bromine  and  their  salts.  No  oxides  of  bromine 
are  known  with  certainty.  The  following  oxygen  acids  have  been  pre- 
pared in  dilute  solutions:  hypobromous  acid  (HBrO),  bromous  acid 
(HBrO2),  bromic  acid  (HBrOg).  The  hypobromites  and  bromates  are 
very  similar  to  the  corresponding  chlorine  compounds  and  are  pre- 
pared by  the  same  general  methods.  From  these  salts  the  free  acids 
can  be  prepared,  as  in  the  case  of  the  corresponding  chlorine  com- 
pounds. The  hypobromites  are  sometimes  used  as  oxidizing  agents, 
decomposing  into  the  corresponding  bromide  and  oxygen. 

The  oxides  and  oxygen  acids  of  iodine.  The  relation  between  the 
oxides  and  the  oxygen  acids  of  iodine  is  shown  in  the  following  table : 

OXIDES  ACIDS 

fH2O  +  I2O      =2HIO,  hypoiodous 

T  ~    |  iodine  monoxide  (unknown)       |  *    .,  .     .     ,, Jr. 

JLO  •{  >    .    -I      acid ;  known  only  in  the  form  of 

t  (anhydride  of  hypoiodous  acid)  J  . 

I       salts 

/  iodine  pentoxide  \ 

LOri  ,./_.•,      ,  .    -,.        .  ^  }•    .     .     .       H,O  +  LO-     =  2  HIOo,  iodic  acid 
2    5  t  (anhydride  of  iodic  acid)  J 

H2O  +  I2O7     =  2  HIO4,  periodic 

acid ;  known  only  in  dilute  solu- 
J  iodine  heptoxide  (unknown)  \ 

'^-'7     1      /•  l^T      •    1  _       _  P  _       •         1  •   _       -        *1X      I 


2    7     (anhydride  of  periodic  acid) 


tion  and  in  the  form  of  salts 
5  H2O  +  I2O7  =  2  H5IO6,  also  known 
as  periodic  acid 

Iodine  pentoxide  (I20B).  This  is  a  white  solid  formed  by  heating  iodic 
acidto200°:  2  HIO8  ^+ H2O  +  I2O6 

The  reaction  .is  reversible,  the  oxide  combining  with  water  at  ordi- 
nary temperatures  to  form  iodic  acid.  At  high  temperatures  it  is 
decomposed  into  its  constituent  elements  and  reacts  therefore  as  an 
oxidizing  agent. 

Iodic  acid  (HI03).  This  acid  is  formed  by  the  action  of  sulfuric  acid 
upon  the  iodates,  although  it  is  more  convenient  to  prepare  it  by 
oxidizing  iodine  directly  with  nitric  acid.  It  forms  white  crystals 
and  is  a  strong  oxidizing  agent. 


274  GENERAL  CHEMISTRY 

Hypoiodites  and  iodates.  These  compounds  are  similar  to  the  cor- 
responding chlorine  compounds  in  their  properties,  and  are  prepared  by 
the  same  general  methods.  Sodium  iodate  is  found  in  Chile  saltpeter. 

Periodic  acid  and  the  periodates.  As  indicated  in  the  table,  the  same 
anhydride  may  combine  with  different  weights  of  water  to  form  dif- 
ferent acids.  The  acids  formed  from  the  same  anhydride  all  belong  to 
the  same  general  class.  Thus,  all  acids  formed  from  the  anhydride 
I  O  are  known  as  periodic  acids.  In.  order  to  distinguish  such  acids 
from  each  other,  certain  prefixes  are  used.  In  the  case  of  the  periodic 
acids,  however,  since  only  one  is  definitely  known  in  the  free  state, 
namely,  H5IO6,  the  term  periodic  acid  is  used  to  designate  this  partic- 
ular compound.  The  periodates  may  be  prepared  by  the  oxidation  of 
the  iodates.  The  one  most  readily  obtained  is  an  acid  salt  and  has  the 
formula  Na2H3IOg.  Periodic  acid  itself,  in  the  form  of  a  white  solid, 
can  be  prepared  from  this.  When  heated  it  decomposes,  forming 
water,  oxygen,  and  iodine  pentoxide. 


CHAPTER  XXI 

CARBON  AND  ITS  COMPOUNDS 

Occurrence.  In  the  free  condition  carbon  is  found  in  nature  in  sev- 
eral forms.  The  diamond  is  practically  pure  carbon.  Coal  and  graphite 
contain  small  percentages  of  other  substances  besides  carbon,  especially 
mineral  matter.  Its  natural  compounds  are  exceedingly  numerous  and 
occur  in  the  form  of  gases,  liquids,  and  solidsi  Carbon  dioxide  is  its 
most  familiar  gaseous  compound.  Natural  gas  and  petroleum  are  com- 
posed principally  of  compounds  of  carbon  and  hydrogen.  The  carbon- 
ates, especially  calcium  carbonate,  constitute  great  strata  of  rocks  and 
are  found  in  almost  every  locality.  Living  organisms,  both  plant  and 
animal,  contain  a  large  percentage  of  combined  carbon,  and  the  number 
of  its  compounds  which  go  to  make  up  all  the  vast  variety  of  animate 
nature  is  almost  limitless.  It  is  commonly  regarded  as  the  element 
most  closely  related  to  life  itself,  although  it  is  undoubtedly  true  that 
the  other  elements  normally  present  in  the  tissues  of  living  organisms 
all  play  an  essential  part  in  the  growth  of  the  organism. 

Forms  of  carbon.  Carbon  occurs  in  a  number  of  different  forms. 
For  purposes  of  study  it  is  convenient  to  divide  these  into  two  general 
classes,  namely,  the  crystalline  and  the  amorphous. 

Crystalline  forms  of  carbon.  Two  forms  of  crystalline  carbon  occur 
in  nature,  namely,  the  diamond  and  graphite. 

1.  Diamond.  This  form  of  carbon  has  long  been  known  and  highly 
prized  as  a  gem.  Diamonds  are  found  in  several  localities,  especially 
in  South  Africa,  the  East  Indies,  and  Brazil.  The  crystals  belong  to 
the  regular  system,  although  the  natural  crystals  are  always  more 
or  less  imperfect.  As  commonly  found  in  nature,  they  are  covered 
with  a  rough  coating.  In  order  to  bring  out  the  brilliancy  of  the 
gem,  the  natural  crystal  is  cut  in  such  a  way  that  the  light  is  most 
effectively  refracted. 

A  pure  diamond  is  perfectly  transparent  and  colorless,  but  many  are  tinted  a 
variety  of  colors  by  traces  of  foreign  substances.  Usually  the  colorless  forms  are 
the  most  highly  prized,  although  in  some  instances  the  color  adds  to  the  value 
as  in  the  case  of  the  famous  Hope  diamond,  which  has  a  beautiful  blue  tint. 

275 


276  GENERAL  CHEMISTRY 

The  weight  of  the  diamond  is  commonly  expressed  in  carats.  Each  carat  is 
equal  to  about  0.2  g.  The  word  carat  is  derived  from  a  Greek  word  meaning 
"  the  seed,  or  bean,  of  the  caroj),  or  locust  tree."  The  beans  were  formerly  used 
in  weighing  diamonds. 

The  largest  diamond  known  was  found  in  the  Transvaal  mines  in  1905,  and 
weighed  3025|  carats.  This  was  known  as  the  Cullinan  diamond,  and  was  pre- 
sented to  King  Edward  VII  by  the  Transvaal  government.  It  was  subsequently 
cut  into  nine  large  stones  and  a  number  of  smaller  ones.  The  two  largest  of 
these  weigh  516.5  and  309T3^  carats  and  are  the  largest  cut  diamonds  in  exist- 
ence. Other  famous  diamonds  are  the  Kohinoor  (106^  carats),  the  Nizam  (277 
carats),  the  Victoria  (180  carats),  and  the  Jubilee  (239  carats). 

Composition  and  properties  of  the  diamond.  The  density  of  the  dia- 
mond is  3.5,  and,  though  brittle,  it  is  one  of  the  hardest  of  substances. 
Specimens  are  often  found  in  nature  which  are  identical,  in  composi- 
tion and  properties,  with  the  ordinary  diamond,  except  that  they  are 
black  and  therefore  valueless  as  gems.  Few  chemical  reagents  have 
any  action  upon  the  diamond,  but  when  heated  in  pure  oxygen  or  air 
it  blackens  and  finally  burns,  forming  carbon  dioxide.  Lavoisier  was 
the  first  to  show  that  carbon  dioxide  is  .formed  by  the  combustion  of 
the  diamond  in  pure  oxygen,  thus  proving  that  it  contained  carbon. 
Later  (1814)  Sir  Humphry  Davy  showed  that  carbon  dioxide  is  the 
sole  product  of  the  combustion,  and  by  determining  the  relation  be- 
tween the  weights  of  the  diamond  burned  and  the  carbon  dioxide 
produced  he  proved  that  the  diamond  is  pure  carbon. 

Artificial  preparation  of  diamonds.  Many  attempts  have  been  made  to  produce 
diamonds  artificially.  These  attempts  were  unsuccessful  until  Moissan,  in  1893, 
finally  succeeded  in  producing  diamonds  identical  in  every  way  with  the  natural 
gem.  The  method  used  by  Moissan  consisted  in  dissolving  pure  carbon  in  molten 
iron  and  quickly  cooling  the  resulting  solution  by  plunging  the  crucible  contain- 
ing it  into  water.  Under  these  conditions  a  portion  of  the  carbon  separated  in 
the  form  of  crystals.  The  iron  was  then  removed  by  dissolving  it  in  acids.  The 
largest  of  the  crystals  so  obtained,  however,  had  a  diameter  of  only  0.5  mm.,  and 
were  thus  too  small  to  have  any  value  as  gems. 

Graphite.  This  form  of  carbon  is  found  in  large  quantities  in  nature, 
especially  in  Ceylon,  Siberia,  and  some  localities  in  the  United  States 
and  Canada.  It  is  a  shining  black  substance,  very  soft  and  greasy  to  the 
touch.  Its  density  is  about  2.3.  It  varies  somewhat  in  properties,  accord- 
ing to  the  locality  in  which  it  is  found.  When  any  form  of  carbon  is 
heated  in  an  electric  furnace  to  a  temperature  of  about  3500°,  it  rapidly 
vaporizes,  and  the  vapor  always  condenses  in  the  form  of  graphite.  This 
property  has  led  to  the  production  of  graphite  on  a  commercial  scale. 


CARBOK  AND  ITS  COMPOUNDS 


277 


The  commercial  production  of  graphite.  The  method  of  producing  graphite  com- 
mercially was  worked  out  by  Acheson.  The  process  consists  essentially  in  heat- 
ing carbon  in  large  electric  furnaces  about  40  ft.  in  length,  a  longitudinal  section 
of  which  is  shown  in  Fig.  99.  The  electrodes  A,  A  are  made  of  graphite.  The 
walls  of  the  furnace  are  built  of  carborundum  and  concrete.  The  furnace  is 
nearly  filled  with  some  form  of  carbon,  such  as  coarse  grains  of  anthracite  coal 
.  Since  anthracite  coal  is  a  poor  conductor  of  electricity,  there  is  placed  in 


FIG.  99 

the  center  of  the  charge  a  core  (C)  of  granulated  carbon,  connecting  the  two 
electrodes,  the  core  serving  to  conduct  the  current  through  the  charge.  The 
charge  is  covered  with  a  mixture  of  sand  and  carbon  (D)  or  similar  materials, 
which  serves  to  exclude  the  air.  An  alternating  current  (40,000  amperes  at  200 
volts)  is  supplied  by  the  generator  G.  Under  the  influence  of  the  intense  heat 
produced  by  the  current  the  carbon  is  changed  into  the  form  of  graphite.  Pre- 
pared in  this  way,  the  product  is  uniform  in  composition  and  free  from  grit, 
and  is  therefore  superior  to  the  natural  product  for  many  purposes. 

Graphite  is  used  in  the  manufacture  of  crucibles,  as  a  lubricant, 
and  as  a  protective  cover  for  iron  in  the  form  of  a  paint  or  polish,  such 
as  stove  polish.  It  has  long  been  used  in  the  manufacture  of  lead 
pencils,  a  fact  which  is  indicated  by  its  name,  which  is  derived  from  a 
Greek  word  meaning  "  to  write." 

Amorphous  carbon.  Many  varieties  of  amorphous  carbon  are  known. 
Some  of  these,  as  the  various  forms  of  coal,  are  found  in  nature, 
while  others,  such  as  charcoal  and  coke,  are  easily  prepared.  These 
forms  differ  merely  in  their  degree  of  purity  and  in  their  physical 
condition.  They  are  of  the  greatest  importance,  owing  to  their  many 
uses  in  the  arts  and  industries. 

1.  Pure  carbon.  Pure  amorphous  carbon  is  best  prepared  by  char- 
ring ordinary  sugar.  This  compound  has  the  composition  expressed 
by  the  formula  C12H22O11.  When  strongly  heated,  the  oxygen  and 
hydrogen  are  expelled  largely  in  the  form  of  water,  while  pure  carbon 
is  left.  Prepared  in  this  way,  carbon  is  a  soft,  very  bulky  black  powder. 
It  was  in  this  way  that  Moissan  obtained  the  carbon  which  he  used 
in  the  preparation  of  artificial  diamonds. 


278 


GENERAL  CHEMISTRY 


2.  Coal  and  coke.  Coals  of  various  kinds  have  been  formed  from 
vast  accumulations  of  vegetable  matter,  which  became  covered  with 
water  and  earthy  material  and  were  thus  protected  from  rapid  decay. 
Under  the   influences  exerted   by  various   geological    agencies    this 
organic  matter  was  slowly  changed  into  coal.    In  anthracite  coal  these 
changes  have  gone  the  farthest,  and  the  carbon  in  this  kind  of  coal  is 
largely  in  the  free  condition.    Soft,  or  bituminous,  coals,  on  the  other 
hand,  contain  a  much  larger  percentage  of  combined  carbon.    When 
heated  strongly  out  of  contact  with  air,  as  in  the  manufacture  of  coal 
gas,  the  carbon  compounds  undergo  complicated  changes  resulting 
in  the  formation  of  a  large  number  of  substances  which  are  given  off 
in  the  form  of  gases  and  vapors,  while  the  mineral  matter  and  free 
carbon  remain  behind  and  constitute  ordinary  coke. 

3.  Charcoal.   This  form  of  carbon  has  long  been  used  as  a  fuel  and  as 
a  reducing  agent  in  obtaining  metals  from  their  oxides.    It  is  prepared 
by  heating  wood  in  the  absence  of  air,  just  as  coke  is  prepared  from  coal 
under  like  conditions.  Formerly  this  process  was  carried  out  in  a  waste- 
ful way  by  merely  covering  piles  of  wood  with  sod  and  then  igniting 
the  wood.    By  this  process  some  of  the  wood  is  burned,  while  the  re- 
mainder is  decomposed  by  the  heat,  forming  charcoal.  In  heating  wood, 
just  as  in  heating  coal,  many  valuable  volatile  products  are  formed. 
In  the  preparation  of  charcoal  by  the  older  methods  these  products  are 
all  lost.   At  present  an  increasing  quantity  of  charcoal  is  made  by  heat- 
ing the  wood  in  large  retorts,  and  the  volatile  products  are  condensed 
and  saved  as  in  the  case  of  coal.   Among  the  products  so  obtained  are 
wood  alcohol,  acetic  acid,  and  acetone,  as  well  as  a  number  of  less  im- 
portant substances.  The  mineral  constituents  of  the  wood  remain  in  the 
charcoal.   The  relative  composition  of  coal,  coke,  wood,  and  charcoal  is 
represented  approximately  by  the  following  analyses  of  typical  samples : 


TOTAL 
CARBON 

HYDRO- 
GEN 

OXYGEN 

NITRO- 
GEN 

SULFUR 

ASH 

TOTAL 
VOLATILE 
MATTER 

Coal  (anthracite)  . 
Coal  (semi- 
bituminous)  .     . 
Coal  (bituminous) 
Wood  . 

82.04% 

82.71% 
78.03% 
40  00% 

2.70% 

4.43% 
4.99% 

7  20% 

3.50% 

3.98% 
6.11% 
50  70% 

0.77% 

1.33% 

1.53% 
0  80*7 

0.74% 

0.68% 

1.05% 

10.25% 

6.87% 
8.29% 
1  30% 

5  to  8% 

18  to  20% 
30  to  35% 
80  to  Q0<7 

Coke  .     . 

89  00% 

traces 

traces 

0  80^ 

10  20*7 

Charcoal. 

97.00% 

traces 

traces 

traces 

traces 

3.00% 

traces 

CAKBON  AND  ITS  COMPOUNDS 


279 


Modern  methods  for  the  production  of  charcoal.  Fig.  100  shows  the  essential 
parts  of  a  modern  plant  for  making  charcoal.  The  iron  cars  A,  A  loaded  with  from 
one  to  three  cords  of  wood  are  run  into  the  retort  B,  which  is  then  made  air- 
tight. The  retort  is  then  heated  slowly  for  about  twenty-four  hours,  the  heat 
being  generated  in  the  fireplace  F,  F.  The  volatile  products  escape  through  the 


pipes  C,  C  and  pass  through  the  condensers  D,  D.  Here  those  portions  which 
are  liquid  at  ordinary  temperatures  (wood  alcohol,  acetic  acid,  and  acetone)  are 
condensed  and  flow  off  through  pipes  E,  E,  to  suitable  containers,  while  the 
gaseous  products  are  led  back  into  the  fireplace  and  burned.  After  the  volatile 
matter  is  expelled  from  the  wood,  the  retort  is  allowed  to  cool  somewhat,  and 
the  cars  containing  the  charcoal  are  run  out  of  the  retort  into  cooling  chambers, 
their  places  in  the  retort  being  taken  by  other  cars  loaded  with  wood. 

4.  Bone  black.  This  form  of  carbon  is  sometimes  called  animal  char- 
coal and  is  made  by  heating  bones  and  animal  refuse  in  the  absence  of 
ah*.  Bones  are  composed  of  about  40  per  cent  organic  matter  and  60 
per  cent  mineral  matter,  chiefly  calcium  phosphate.  When  heated  in 
the  absence  of  air  the  organic  matter  is  decomposed,  resulting  in  the 
formation  of  volatile  matter  and  free  carbon,  which  remains,  in  a  finely 
divided  state,  scattered  through  the  mineral  portion  of  the  bone.  The 
bone  black  so  obtained  consists  principally  of  calcium  phosphate  with 
a  relatively  small  percentage  of  carbon.  For  some  uses  it  is  desirable 
that  the  mineral  part  be  removed,  and  this  is  done  by  the  action  of 
hydrochloric  acid,  which  dissolves  the  calcium  phosphate  present  but 
has  no  action  upon  the  carbon.  For  most  purposes,  however,  the  pres- 
ence of  the  calcium  phosphate  is  not  objectionable. 


280  GENERAL  CHEMISTRY 

The  volatile  matter  formed  in  the  heating  of  bones  condenses  to  a  dark-colored, 
foul-smelling  liquid  known  as  bone  oil.  It  is  a  mixture  of  a  number  of  compounds, 
important  among  which  is  pyridine  (C5H5N),  a  compound  sometimes  added  to 
alcohol  to  render  it  unfit  for  drinking  (p.  303). 

Destructive  distillation.  The  process  of  decomposing  such  substances 
as  coal,  wood,  and  bones  by  heating  them  in  the  absence  of  air  is 
termed  destructive  distillation.  As  commonly  expressed,  coke,  char- 
coal, and  bone  black  are  made  by  the  destructive  distillation  of  coal, 
wood,  and  bones,  respectively. 

5.  Lampblack  (soot).  If  a  piece  of  cold  porcelain  is  held  for  a  few 
seconds  in  the  flame  of  a  candle,  the  temperature  of  the  flame  is  reduced 
to  such  an  extent  that  much  of  the  carbon  present  no  longer  burns 
but  is  deposited  on  the  porcelain  in  the  form  of  a  black  material  known 
as  lampblack,  or  soot.  It  is  manufactured  on  a  large  scale  by  methods 
based  on  this  same  general  principle,  and  is  used  for  various  purposes, 
especially  in  the  manufacture  of  printer's  ink. 

Properties  of  carbon.  The  various  forms  of  carbon  are  all  odorless, 
tasteless  solids.  They  differ,  however,  in  many  properties,  especially  in 
color,  density,  and  hardness.  Carbon  is  insoluble  in  all  ordinary  sol- 
vents. Some  of  the  metals,  such  as  iron,  gold,  and  silver,  when  melted, 
dissolve  it,  forming  a  solution  from  which  the  carbon  can  be  separated 
unchanged.  Melted  iron  is  the  best  solvent,  dissolving  about  1  per  cent 
of  its  weight  of  carbon.  In  the  form  of  bone  black  or  charcoal,  carbon 
has  the  property  of  absorbing  relatively  large  quantities  of  certain 
gases.  For  example,  1  volume  of  charcoal  absorbs  about  178  volumes 
of  ammonia  and  166  volumes  of  hydrogen  sulfide,  the  exact  volume 
of  the  gas  absorbed  depending  upon  the  physical  condition  of  the 
carbon  as  well  as  upon  the  temperature  and  the  pressure.  Similarly, 
it  absorbs  certain  kinds  of  organic  matter  from  their  solutions.  Thus, 
water  colored  with  litmus,  when  heated  with  bone  black  and  filtered, 
is  entirely  decolorized. 

Carbon  is  characterized  by  its  great  stability  toward  heat,  but  the 
fact  that  a  thin  film  of  the  element  collects  on  the  interior  surface  of 
electric-light  bulbs  after  continued  usage  shows  that  it  has  a  percep- 
tible vapor  pressure  at  the  temperature  reached  in  the  incandescent 
lamp.  At  the  temperature  of  the  electric  arc  (about  3500°)  the  vapor 
pressure  of  carbon  is  greater  than  the  atmospheric  pressure,  so  that  at 
this  temperature  the  element  rapidly  vaporizes,  passing  directly  from 
the  solid  into  the  gaseous  state. 


CARBON  AND  ITS  COMPOUNDS          281 

Adsorption.  When  carbon  is  brought  in  contact  with  a  gas  or  with  solutions  of 
certain  compounds,  especially  the  more  complex  constituents  of  organic  matter, 
the  molecules  of  the  gas  or  of  the  dissolved  compound  apparently  condense  upon 
the  surface  of  the  carbon.  This  phenomenon  is  known  as  adsorption.  Many  other 
substances  besides  carbon  act  in  a  similar  way.  Carbon,  however,  is  one  of  the 
most  efficient  adsorbing  agents,  owing  largely  to  the  fact  that  it  is  very  porous 
and  hence  presents  a  comparatively  large  condensing  surface. 

Chemical  conduct.  At  ordinary  temperatures  carbon  is  a  very  inert 
substance,  but  at  higher  temperatures  it  combines  directly  with  a  num- 
ber of  elements,  such  as  oxygen,  hydrogen,  sulfur,  nitrogen,  silicon, 
boron,  and  the  halogens.  Because  of  its  strong  affinity  for  oxygen  it  is 
an  excellent  reducing  agent.  Carbon  also  combines  directly  with  many 
of  the  metals,  forming  compounds  called  carbides.  One  of  the  most 
important  of  these  is  calcium  carbide  (CaC2),  used  so  largely  in  the 
preparation  of  acetylene.  When  heated  in  the  presence  of  oxygen, 
carbon  burns,  forming  carbon  dioxide. 

Uses  of  carbon.  The  chief  use  of  amorphous  carbon  is  for  fuel,  to 
furnish  heat  and  power  for  all  the  uses  of  civilization.  An  enormous 
quantity  of  carbon,  in  the  form  of  coal,  coke,  and  charcoal,  is  used  as  a 
reducing  agent  in  the  separation  of  the  various  metals  from  their  ores. 
Lampblack  is  used  for  making  indelible  ink,  printer's  ink,  and  black 
varnishes,  while  bone  black  and  charcoal  are  used  in  water  niters.  In 
the  refining  of  sugar  the  dark  solution  of  the  impure  compound  is  fil- 
tered through  layers  of  bone  black,  which  removes  the  coloring  matter. 
On  evaporation  the  resulting  solution  yields  the  colorless  sugar. 

Compounds  of  carbon.  The  compounds  of  carbon  are  more  numerous 
by  far  than  are  the  compounds  of  any  other  element.  Nearly  200,000 
of  them  have  been  described,  and  additional  ones  are  being  continually 
added  to  the  list.  The  existence  of  such  a  large  number  of  compounds 
is  due  to  the  property  which  the  carbon  atoms  possess  of  combining 
with  each  other  and  thus  building  up  compounds  more  or  less  complex 
in  character.  Because  of  the  large  number  of  the  compounds  of  carbon, 
and  also  because  of  certain  well-defined  characteristics  which  these 
compounds  possess,  it  has  been  found  convenient  to  include  them 
in  a  separate  course  of  study,  which  is  known  as  the  chemistry  of 
the  compounds  of  carbon,  or,  more  commonly,  as  organic  chemistry. 

The  selection  of  the  term  organic  chemistry  dates  back  to  an  early  period,  when 
it  included  simply  those  compounds  of  carbon  found  in  living  organisms.  It  was 
supposed  that  these  compounds  could  only  be  formed  through  the  influence  of  the 
living  or  vital  force  of  the  organisms,  and  hence  that  it  was  impossible  to  prepare 


282  GENERAL  CHEMISTRY 

them  in  the  laboratory  by  synthetic  methods.  Finally,  in  1828,  the  German  chemist 
Wohler  prepared  urea,  a  typical  organic  compound,  by  synthetic  methods,  thus 
showing  that  the  existing  conception  was  an  erroneous  one.  The  term  organic 
chemistry,  however,  has  been  retained  as  a  convenient  one  for  designating  the 
chemistry  of  all  carbon  compounds.  This  includes  most  of  the  compounds  present 
in  organisms,  these  being,  however,  a  relatively  small  percentage  of  the  total 
number  of  carbon  compounds  known.  It  is  not  advisable  for  us  to  study  any  large 
number  of  these  compounds  at  present.  Only  a  few  of  the  more  important  ones 
will  be  discussed. 

THE  OXIDES  OF  CARBON  AND  CARBONIC   ACID 

Carbon  forms  three  oxides:  namely,  carbon  monoxide  (CO),  carbon 
dioxide  (CO2),  carbon  suboxide  (C3O2).  They  are  all  colorless  gases. 
But  little  is  known  of  the  suboxide,  and  no  further  mention  will  be 
made  of  it.  Carbon  dioxide,  being  the  most  abundant  and  the  best 
known,  will  be  discussed  first. 

Carbon  dioxide  (carbonic  anhydride)  (C02).  This  compound  is  pres- 
ent in  the  open  air  to  the  extent  of  from  3  to  4  parts  -in  10,000, 
and  this  apparently  small  percentage  is  of  fundamental  importance  in 
nature.  In  some  localities  it  escapes  from  the  earth  in  great  quantities, 
and  many  spring  waters  contain  it  in  solution.  When  such  waters 
reach  the  surface  of  the  earth,  the  pressure  upon  them  is  diminished 
and  the  gas  escapes  with  effervescence.  Carbon  dioxide  is  a  product 
of  the  oxidation  of  all  organic  matter,  and  is  therefore  formed  in  the 
process  of  combustion,  as  well  as  in  that  of  decay.  It  is  exhaled  from 
the  lungs  of  all  animals  in  respiration,  and  is  a  product  of  many  fer- 
mentation processes,  such  as  that  which  takes  place  in  the  manufacture 
of  alcoholic  liquors. 

Preparation.  In  the  laboratory  carbon  dioxide  is  prepared -by  the 
action  of  an  acid,  such  as  hydrochloric  or  sulfuric,  upon  some  salt  of 
carbonic  acid  (H2COg).  These  salts  are  termed  carbonates.  The  car- 
bonate generally  used  is  that  of  calcium  (CaCOg),  which  occurs 
abundantly  in  nature  in  the  form  of  limestone  and  marble.  When 
hydrochloric  or  sulfuric  acid  is  added  to  a  carbonate,  carbonic  acid  is 
formed,  just  as  one  would  expect  (p.  200).  This  acid,  however,  as  fast 
as  formed,  decomposes  into  water  and  carbon  dioxide.  The  latter,  being 
but  moderately  soluble,  escapes  and  may  be  collected  by  displacement 
of  air  or  water.  The  equations  for  the  reaction  are  as  follows : 

CaCO,  +  2  HC1  =  CaCl2  +  H2COg 
H9C03  =  H0  +  C0 


CARBON  AND  ITS  COMPOUNDS  283 

To  prepare  the  gas  in  the  laboratory,  pieces  of  marble  are  placed  in  the  gen- 
erator A  (Fig.  81,  p.  208),  and  commercial  hydrochloric  acid,  diluted  with  an  equal 
volume  of  water,  is  added  slowly  from  the  separatory  funnel  B.  The  gas  escapes 
through  C  and  may  be  collected  in  cylinders.  The  Kipp  apparatus  (Fig.  18)  is 
much  more  convenient. 

Properties.  Carbon  dioxide  is  a  colorless,  practically  odorless  gas 
1.5  times  as  heavy  as  air.  Its  weight  may  be  inferred  from  the  fact 
that  it  can  be  siphoned  or  poured  like  water  from  one  vessel  down- 
ward into  another.  At  15°,  and  under  ordinary  pressure,  1  volume 
of  water  dissolves  1  volume  of  the  gas.  The  resulting  solution  has 
a  somewhat  biting,  pungent  taste.  At  ordinary  temperatures  <(20°) 
carbon  dioxide  is  liquefied  by  a  pressure  of  56.3  atmospheres.  Liquid 
carbon  dioxide  is  colorless  and  slightly  lighter  than  water.  It  dissolves 
some  organic  substances,  such  as  naphthalene  and  camphor,  but  only 
a  very  few  inorganic  substances.  The  commercial  carbon  dioxide, 
compressed  in  steel  cylinders,  is  under  such  great  pressure  that  it  is 
largely  in  the  liquid  state.  When  the  pressure  is  removed,  the  rapid 
expansion  of  the  gas  reduces  the  temperature  sufficiently  to  freeze  a 
portion  of  the  escaping  liquid  to  a  snowlike  solid. 

It  is  a  very  simple  matter  to  obtain  this  solid  carbon  dioxide  and  to  show  its 
low  temperature  by  freezing  mercury  with  it.  Iron  cylinders  filled  with  carbon 
dioxide  under  pressure  are  inexpensive  and  easily  available.  To  obtain  the  solid 
carbon  dioxide  the  cylinder  should  be  placed  across  a  desk  and  supported  in  such  a 
way  that  the  end  provided  with  a  stopcock  is  several  inches  lower  than  the  other 
end.  A  loose  bag  is  made  by  holding  the  corners  of  a  piece  of  cloth  tightly 
around  the  neck  of  the  stopcock.  Upon  opening  the  stopcock  the  liquid,  together 
with  the  gas  formed  by  its  rapid  evaporation,  rushes  out.  The  heat  absorbed 
by  the  evaporation  freezes  a  portion  of  the  liquid,  which  is  strained  out  from  the 
gas  by  the  cloth  bag.  A  considerable  quantity  of  the  snow  very  soon  collects 
in  the  bag.  Mercury  may  be  frozen  by  this  snow  in  the  following  way :  A  filter 
paper  is  placed  in  the  bottom  of  a  small  evaporating  dish  and  some  mercury 
poured  upon  it.  One  end  of  a  piece  of  wire  is  wound  into  a  flat  coil  and  dipped 
into  the  mercury.  A  quantity  of  the  solid  carbon  dioxide  is  placed  upon  the  mer- 
cury and  from  10  to  15  cc.  of  ether  poured  over  it.  The  temperature  is  reduced 
to-—  50°,  so  that  the  mercury  solidifies  in  a  minute  or  two  and  may  be  removed 
from  the  dish  by  the  wire  which  serves  as  a  handle.  While  the  solid  is  intensely 
cold,  it  may  be  handled  without  danger,  because  the  skin  is  protected  from  direct 
contact  with  it  by  a  layer  of  gas.  The  ether  is  added  to  the  snow  in  freezing 
mercury,  since  in  this  way  better  contact  is  secured. 

Chemical  conduct.  Carbon  dioxide  is  a  very  stable  substance.  At 
high  temperatures  partial  decomposition  takes  place,  as  expressed  in 
the  following  equation: 


284 


GENERAL  CHEMISTRY 


At  2000°,  under  a  pressure  of  6  atmospheres,  about  5  per  cent  of  the 
carbon  dioxide  is  thus  decomposed.  It  will  not  combine  with  oxygen 
and  is  therefore  incombustible  ;  neither  will  other  substances  burn  in 
it  under  ordinary  conditions,  for  although  it  contains  a  large  percent- 
age of  oxygen,  this  is  held  in  very  firm  combination.  In  this  respect 
it  differs  from  such  oxides  as  nitrogen  dioxide,  which  readily  support 
combustion.  A  few  energetic  reducing  agents  remove  at  least  a  part 
of  its  oxygen.  Thus,  if  it  is  passed  over  carbon  at  temperatures  above 
1000°,  the  gas  is  partially  reduced,  forming  carbon  monoxide  : 

C0 


At  high  temperatures  sodium  reduces  carbon  dioxide  to  carbon  : 

3  CO2  +  4  Na  =  2  Na2CO3,+  C 

Carbon  dioxide  combines  with  some  of  the  metallic  oxides,  forming 

CO0:«z±CaCO. 


carbonates : 


Uses  of  carbon  dioxide.  Carbon  dioxide  is  obtained  as  a  by-product 
in  a  number  of  chemical  processes,  especially  in  the  preparation  of 
alcoholic  liquors  and  alcohol.  It  is  pumped  into  strong  steel  cylinders, 
and  in  this  form  is  an  article  of  commerce.  It  is  used  chiefly  in  the 
manufacture  of  soda  water  and  similar  beverages,  and  as  a  fire  extin- 
guisher. Ordinary  soda  water  is  simply  water  charged  with  carbon 

dioxide  under  pressure.  When  the  pres- 
sure is  removed,  the  excess  of  gas  escapes, 
producing  effervescence.  Most  of  the  port- 
able fire  extinguishers  are  simply  devices 
for  generating  carbon  dioxide.  It  is  not 
necessary  that  all  the  oxygen  should  be 
kept  away  from  a  fire  in  order  to  smother 
it.  A  burning  candle,  for  example,  is  ex- 
tinguished in  air  which  contains  only  2.5 
per  cent  of  carbon  dioxide. 


The  general  type  of  the  portable  fire  extin- 
guisher is  shown  in  Fig.  101.  The  liquid  is  a 
solution  of  sodium  hydrogen  carbonate  in  water. 
The  bottle  A  contains  sulf  uric  acid  in  sufficient 
amount  to  react  with  the  sodium  carbonate  in 

solution.  In  case  of  fire  the  bottle  containing  the  sulfuric  acid  is  broken  by  forc- 
ing down  the  rod  B.  The  sulfuric  acid  immediately  reacts  with  the  carbonate, 
generating  carbon  dioxide,  some  of  which  dissolves  in  the  water,  while  the 


FIG.  101 


CARBON  AND  ITS  COMPOUNDS          285 

remainder  forces  the  solution  out  through  the  nozzle  C.  While  the  total  quantity 
of  water  furnished  by  such  an  extinguisher  is  comparatively  small,  it  is  very 
effective  as  a  fire  extinguisher,  because  of  the  large  percentage  of  carbon  dioxide 
which  it  contains  in  solution. 

Carbonic  acid  (H2C03).  This  acid  is  unstable  and  is  known  only  in 
the  form  of  a  very  dilute  solution.  This  solution  is  most  readily  pre- 
pared by  passing  carbon  dioxide  into  water : 

H20+C02^^H2CO,  (1) 

The  volume  of  carbon  dioxide  absorbed  in  pure  water  is  relatively 
small.  If,  however,  the  water  contains  a  base,  such  as  sodium  hy- 
droxide, in  solution,  the  carbonic  acid  formed  according  to  equation 
(1)  reacts  with  the  base  to  form  the  corresponding  carbonate: 

H2C08  +  2  NaOH  +=+  Na2CO3  +  2  H2O  (2) 

The  removal  of  the  carbonic  acid  results  in  the  union  of  more  carbon 
dioxide  and  water,  according  to  equation  (1),  so  that  the  absorption 
of  carbon  dioxide  will  continue  until  practically  all  of  the  base  has 
been  changed  into  the  corresponding  carbonate. 

The  following  structural  formula  is  in  best  accord  with  the  con- 
duct of  carbonic  acid:  ur. 

HO.  r  _  o 

HO>C 

Salts  of  carbonic  acid ;  the  carbonates.  Since  carbonic  acid  is  a 
dibasic  acid,  it  forms  both  normal  and  acid  salts. 

1.  Normal  carbonates.  The  normal  carbonates  are  found  in  large 
quantities  in  nature  and  are  often  used  in  chemical  processes.  Some 
of  these  are  well-known  compounds.  Thus,  ordinary  limestone  is  a 
more  or  less  impure  form  of  calcium  carbonate.  Marble  is  nearly  pure 
calcium  carbonate  in  crystalline  condition.  Normal  sodium  carbonate 
(Na2CO3)  is  the  well-known  soda  ash,  so  largely  used  in  the  manu- 
facture of  soap  and  glass.  Among  the  normal  carbonates  only  those 
of  sodium,  potassium,  and  ammonium  are  soluble  in  water,  and  these 
can  be  prepared  by  passing  carbon  dioxide  into  solutions  of  the  base, 
as  previously  explained.  The  insoluble  carbonates  can  be  prepared  by 
the  general  method  for  preparing  insoluble  compounds.  Thus,  calcium 
carbonate  is  formed  when  a  solution  of  sodium  carbonate  is  added 
to  a  solution  of  any  compound  of  calcium,  such  as  calcium  chloride : 

Na2CO3  +  CaCl2  =  2  NaCl  +  CaCO8 


286  GENERAL  CHEMISTRY 

Since  carbonic  acid  is  such  a  weak  acid,  and  so  readily  decomposed, 
almost  any  acid  will  act  upon  its  salts  with  corresponding  evolution 
of  carbon  dioxide.  This  reaction  is  used  as  a  test  for  carbonates, 
since  the  carbon  dioxide  evolved  can  readily  be  detected.  Most  of  the 
carbonates  are  decomposed  by  heat.  Ordinary  lime,  for  example,  is 
made  by  strongly  heating  calcium  carbonate  : 

CaCO3:<=±CaO  +  CO2 

2.  Acid  carbonates.  The  acid  carbonates  are  made  by  treating  a 
normal  carbonate  with  an  excess  of  carbonic  acid.  The  most  impor- 
tant of  these  is  sodium  acid  carbonate  (NaHCO3),  or  ordinary  baking 
soda.  With  few  exceptions  they  are  very  unstable  and,  when  heated, 
readily  decompose  even  in  solution.  The  preparation  and  properties 
of  the  acid  carbonates  may  be  illustrated  by  a  single  example.  If 
carbon  dioxide  is  passed  into  a  solution  of  calcium  hydroxide  (lime- 
water),  calcium  carbonate  at  first  precipitates  : 


H2C03  +  Ca(OH)2  =  CaC03  +  2  H2O 

If  the  current  of  carbon  dioxide  is  continued,  however,  the  precipi- 
tated. calcium  carbonate  soon  dissolves.  This  is  due  to  the  formation 
of  calcium  hydrogen  carbonate,  which,  being  soluble,  dissolves  in  the 
water  present:  CaCO3  +  H2CO3  =  Ca(HCO3)2  jjjg. 

If  now  the  solution  is  heated,  the  acid  carbonate  decomposes,  and 
calcium  carbonate  once  more  precipitates: 

Ca(HC03)2  =  CaC03  +  H2O  +  CO2 

Carbon  monoxide  (CO).  Carbon  monoxide  occurs  in  the  gases  issu- 
ing from  volcanoes.  It  can  be  prepared  in  a  number  of  ways,  the  most 
important  of  which  are  the  following  : 

1.  By  the  partial  reduction  of  carbon  dioxide.  When  carbon  dioxide  is 
conducted  over  highly  heated  carbon,  the  monoxide  results  : 


When  coal  burns  in  a  stove,  carbon  dioxide  is  at  first  formed  in  the 
free  supply  of  air,  but  as  the  hot  gas  rises  through  the  glowing  coals 
it  is  reduced  to  carbon  monoxide.  When  this  gas  comes  in  contact 
with  the  air  above  the  coal,  it  combines  with  oxygen  to  form  carbon 


CAKBON  AOT)  ITS  COMPOUNDS  287 

dioxide,  burning  with  the  blue  flame  so  often  noticed  above  a  bed  of 
coals,  especially  in  the  case  of  hard  coal. 

2.  By  the  decomposition  of  oxalic  acid.  In  the  laboratory,  carbon 
monoxide  is  usually  prepared  by  the  action  of  concentrated  sulfuric 
acid  upon  the  hydrate  of  oxalic  acid,  a  compound  having  the  formula 


C2H204-2H20: 


C2H204  •  2  H20  =  3  H20  +  CO2  +  CO 


The  sulfuric  acid  assists  in  the  process  by  absorbing  the  water  as  fast 
as  it  is  formed.  The  resulting  mixture  of  carbon  dioxide  and  carbon 
monoxide  is  made  to  bubble  through  a  solution  of  a  base  such  as 
sodium  hydroxide  or  calcium  hydroxide,  which  combines  with  the 
carbon  dioxide.  The  carbon  monoxide  is  thus  obtained  in  the  pure 
form,  and  may  be  collected  over  water. 

Properties.  Carbon  monoxide  is  a  colorless,  practically  odorless 
gas.  It  is  0.967  times  as  heavy  as  air,  and  is  so  difficult  to  liquefy 
that  it  was  formerly  regarded  as  one  of  the  permanent  gases.  Its 
critical  temperature  is  about  — 141°  and  its  critical  pressure  36  at- 
mospheres. While  almost  insoluble  in  water,  it  is  absorbed  by  a 
number  of  organic  liquids,  and  especially  by  a  solution  of  cuprous 
chloride  containing  either  hydrochloric  acid  or  ammonia.  It  is  a 
very  active  compound,  combining  directly  with  a  great  many  sub- 
stances. It  has  a  marked  affinity  for  oxygen  and  burns  with  a  blue 
flame.  It  is  therefore  a  strong  reducing  agent.  For  example,  when 
it  is  passed  over  copper  oxide  heated  in  a  tube,  the  copper  is  reduced 
to  the  metallic  state : 

CuO  +  CO  =  Cu  +  CO2 

Carbon  monoxide  also  combines  with  chlorine,  sulfur,  and  some  of 
the  metals,  such  as  nickel  and  iron.  It  is  very  poisonous  when  in- 
haled. Deaths  not  infrequently  result  from  the  stoppage  of  stove- 
pipes or  chimneys.  The  draft  of  air  is  thereby  diminished  to  such  an 
extent  that  carbon  monoxide,  rather  than  dioxide,  forms  and,  not 
having  egress  through  the  chimney,  escapes  into  the  room.  It  is  a 
very  treacherous  poison,  since  it  is  practically  odorless. 

The  reducing  power  of  carbon  monoxide.  Fig.  102  illustrates  a  method  of  showing 
the  reducing  power  of  carbon  monoxide.  The  gas  is  generated  by  gently  heating 
a  mixture  of  oxalic  acid  and  sulfuric  acid  in  the  flask  A.  The  bottle  B  contains  a 
solution  of  sodium  hydroxide,  which  removes  the  carbon  dioxide  formed  along 
with  the  monoxide.  C  contains  a  solution  of  calcium  hydroxide,  which  serves  to 
show  that  all  the  carbon  dioxide  has  been  removed,  since  its  presence  in  the  gas 


288 


GENERAL   CHEMISTRY 


would  cause  a  precipitate  of  calcium  carbonate.  E  is  a  hard-glass  tube  contain- 
ing copper  oxide,  which  is  heated  by  a  burner.  The  black  copper  oxide  is  reduced 
to  reddish  metallic  copper  by  the  carbon  monoxide,  which  is  thereby  changed  to 
carbon  dioxide.  The  presence  of  the  carbon  dioxide  is  shown  by  the  precipitate 
in  the  calcium  hydroxide  solution  in  D.  Any  unchanged  carbon  monoxide  is 
collected  over  water  in  F. 


FIG.  102 

Structural  formulas  of  the  oxides  of  carbon.  The  structural  formulas 
of  carbon  dioxide  and  carbon  monoxide  are  as  follows : 


=  0  = 


c  =  o 


It  will  be  noted  that  in  carbon  dioxide  the  carbon  is  tetravalent,  while 
in  carbon  monoxide  it  is  only  divalent.  In  the  formation  of  its  com- 
pounds carbon  is  normally  tetravalent.  In  the  relatively  few  known 
instances  in  which  it  has  a  lower  valence,  as  in  carbon  monoxide,  the 
compound  shows  a  marked  tendency  to  combine  with  a  divalent  group 
or  element  or  two  univalent  groups  or  elements,  the  carbon  thus  pass- 
ing to  the  normal  tetravalent  condition.  Those  compounds,  therefore, 
in  which  the  carbon  has  a  valence  of  less  than  4  are  always  very 
reactive. 

Some  simple  derivatives  of  carbonic  acid.  The  structural  relation  of 
carbonic  acid  to  three  of  its  important  derivatives  can  be  seen  by  com- 
paring their  structural  formulas : 


HO 
HO 


>C=O 


>C  = 


NH 
HO 


2>C  = 


NH 
NH 


(carbonic  acid) 


(carbonyl  chloride) 


(carbamic  acid) 


(urea) 


In  carbonyl  chloride   each  of  the  hydroxyl  groups  of  carbonic  acid 
has  been  displaced  by  chlorine,  while  in  carbamic  acid  one  hydroxyl 


CARBON  AND  ITS  COMPOUNDS  289 

group,  and  in  urea  both  hydroxyl  groups,  have  been  displaced  by 
the  amido  group  NH2. 

Carbonyl  chloride  (phosgene)  (COC12).  This  compound  is  prepared 
by  passing  a  mixture  of  carbon  monoxide  and  chlorine  over  animal 
charcoal,  which  acts  as  a  catalytic  agent  : 

CO  +  C12  =  COC12 

The  union  of  the  two  gases  is  also  greatly  accelerated  by  sunlight, 
a  fact  indicated  by  the  name  phosgene,  which  means  "  generated  by 
light."  Carbonyl  chloride  is  a  colorless  gas,  easily  condensed  to  a 
liquid  boiling  at  8°.  With  water  it  forms  carbonic  and  hydrochloric 
acids  :  _ 

HO|H_C!]>C=O    =   2Hci  +  ™>c=o 

Carbonyl  chloride  bears  to  carbonic  acid  exactly  the  same  relation 
that  sulfuryl  chloride  bears  to  sulfuric  acid  (p.  230).  They  both  be- 
long to  the  general  group  known  as  the  chlorides  of  acids. 

Urea  (CO(NH2)2).  Urea  is  formed  by  the  action  of  carbonyl  chlo- 
ride upon  ammonia: 


It  is  a  white,  crystalline  solid,  very  soluble  in  water.  .  Most  of 
the  waste  nitrogenous  matter  in  the  human  body  is  eliminated  in  the 
liquid  excretions  in  the  form  of  urea.  Oxidizing  agents  convert  it 
into  water,  carbon  dioxide,  and  nitrogen  : 

2  CO(NH2)2  +  3  02  =  2  N2  +  2  CO2  +  4  H2O 

This  reaction  is  often  used  for  the  approximate  determination  of  the  percent- 
age of  urea  in  urine,  either  sodium  hypochlorite  or  sodium  hypobromite  being 
used  as  the  oxidizing  agent.  The  evolved  gases  are  bubbled  through  a  solution 
of  sodium  hydroxide,  which  absorbs  the  carbon  dioxide.  The  quantity  of  urea 
present  is  calculated  from  the  volume  of  the  resulting  nitrogen. 

CARBON  BISULFIDE  AND  THIOCARBONIC  ACID 

These  compounds  differ  from  carbon  dioxide  and  carbonic  acid  in 
composition  in  that  they  contain  sulfur  in  place  of  oxygen. 

Carbon  disulfide  (CS2).  When  sulfur.  vapor  is  passed  over  highly 
heated  carbon,  the  two  elements  combine,  forming  carbon  disulfide, 
just  as  carbon  and  oxygen  unite  to  form  carbon  dioxide.  Carbon  disul- 
fide is  a  heavy,  colorless,  highly  refractive  liquid  which  boils  at  46°. 


290 


GENERAL  CHEMISTRY 


When  pure  it  has  a  pleasant  odor,  but  it  gradually  undergoes  slight 
decomposition  and  acquires  a  most  disagreeable  odor.  When  passed 
through  heated  tubes  it  decomposes : 


CS, 


s 


Its  vapor  is  very  inflammable,  burning  in  the  air  to  form  carbon 
dioxide  and  sulfur  dioxide: 


Carbon  disulfide  is  a  good  solvent  for  many  substances,  such  as  gums, 
resins,  and  waxes,  which  are  not  soluble  in  most  liquids,  and  it  is  there- 
fore used  as  a  solvent  for  such  substances.  It  is  also  used  as  an  insec- 
ticide. Its  vapor  is  poisonous  as  well  as  highly  inflammable,  so  that 
one  must  exercise  great  care  in  working  with  it. 

• 
Commercial   preparation  of   carbon  disulfide.    The  process  used  at  present  in 

the  manufacture  of  carbon  disulfide  consists  essentially  in  passing  an  electric 
current  through  a  mixture  of  charcoal  and  sulfur 
under  such  conditions  that  the  heat  generated  by  the 
current  is  sufficient  to  raise  the  temperature  of  the 
mixture  to  a  point  at  which  the  carbon  and  sulfur 
combine.  Fig.  103  represents  a  section  of  a  type  of 
furnace  devised  by  Taylor  for  heating  the  mixture. 
The  furnace  is  filled  with  charcoal  A,  supplied  from 
the  hopper  B  by  lowering  the  metal  cone  C.  Sulfur 
is  admitted  from  the  hoppers  D,  D.  The  wires  from 
the  dynamo  lead  in  through  the  openings  at  E,  E. 
The  connections  are  so  made  that  carbon  rods  led  in 
through  the  tubes  F,  F  conduct  the  current  toward 
the  bottom  and  center  of  the  furnace.  Here  the  heat 
generated  is  such  that  the  carbon  and  sulfur  combine. 
The  resulting  vapors  of  carbon  disulfide  pass  up 
through  the  furnace,  escaping  through  the  tube  H, 
from  which  they  are  conducted  to  a  suitable  condens- 
ing apparatus.  The  largest  of  these  furnaces  are  16  ft. 
in  diameter  and  41  ft.  in  height,  and  yield  as  much  as 
FIG.  103  25,000  Ib.  of  the  disulfide  in  24  hours. 

Thiocarbonic  acid.  Corresponding  to  carbonic  acid  and^the  carbonates 
we  have  thiocarbonic  acid  (H2CS3)  and  the  thiocarbonates,  such  as 
CaCS8.  Likewise,  corresponding  to  carbamic  acid  and  urea,  we  have 
thiocarbamic  acid  and  thiourea.  The  reactions  which  these  compounds 
undergo  are  in  general  quite  similar  to  those  of  the  corresponding 
oxygen  compounds. 


CARBON  AND  ITS  COMPOUNDS  291 

CYANOGEN  AND  SOME  RELATED  COMPOUNDS 

Cyanogen  (CN)2.  When  electric  sparks  are  passed  between  carbon 
poles  surrounded  by  nitrogen,  some  of  the  carbon  and  nitrogen  unite 
to  form  a  colorless  gas  known  as  cyanogen,  which  has  the  formula 
(CN)2  or  C2N2.  The  compound  is  much  more  readily  prepared  by 
heating  mercuric  cyanide  (Hg(NC)2).  The  decomposition  of  the 
cyanide  by  heat  is  very  similar  to  the  decomposition  of  the  oxide,  as 
represented  in  the  following  equations : 

2HgO  =  ZHg.+  0, 

Hg(NC)2=Hg  +  C2N2 

Cyanogen  is  a  colorless  gas  with  an  odor  somewhat  like  that  of  peach 
kernels.  It  is  extremely  poisonous.  It  burns  readily,  forming  carbon 
dioxide  and  nitrogen.  At  high  temperatures  it  combines  with  potas- 
sium to  form  potassium  cyanide  (KNC),  which  is  the  potassium  salt 
of  hydrocyanic  acid  (HNC). 

Hydrogen  cyanide  (HNC).  This  compound,  first  obtained  by  Scheele, 
is  well  known  because  of  its  intensely  poisonous  properties.  It  can  be 
prepared  by  the  action  of  sulfuric  acid  upon  the  metallic  cyanides : 

KNC  +  H2SO4  =  HNC  +  KHSO4 

It  is  a  light  colorless  liquid  boiling  at  26.1°.  Its  odor  is  like  that 
of  peach  kernels  or  oil  of  bitter  almonds.  It  mixes  with  water  in  all 
proportions,  forming  the  solution  known  as  hydrocyanic  acid,  or,  more 
commonly,  as  prussic  acid.  It  is  one  of  the  weakest  of  all  acids,  so  that 
its  poisonous  action  is  not  due  to  its  acid  properties.  Its  salts  are 
called  cyanides.  Potassium  cyanide  (KNC)  and  sodium  cyanide 
(NaNC)  are  the  best  known.  They  are  white  solids  and  extremely 
poisonous.  Their  solutions  in  water  react  basic  (p.  225).  As  with 
carbon  monoxide,  the  carbon  present  in  the  cyanides  is  probably 
divalent,  as  shown  in  the  following  structural  formulas: 

H-N  =  C  K-N  =  C 

Cyanic  acid  and  the  cyanates.  When  potassium  cyanide  is  heated 
with  an  oxidizing  agent,  the  white  solid  known  as  potassium  cyanate 
is  formed:  KNC  +  O  =  KNCO 

This  is  the  potassium  salt  of  the  unstable  cyanic  acid  HNCO.  Ammo- 
nium cyanate  has  the  composition  NH^NCO.  When  this  is  dissolved 
in  water  and  the  solution  heated,  urea  is  formed.  It  was  by  this 


292  GENERAL  CHEMISTRY 

method  that  Wohler  synthesized  urea  (p.  289).  The  change  that  takes 
place  when  ammonium  cyanate  is  heated  is  represented  by  the  follow- 
ing equation :  NH 

NH4-N  =  C  =  0    -+    £%>C  =  0 

Corresponding  to  cyanic  acid  and  the  cyanates,  we  have  also  thio- 
cyanic  acid  and  the  thiocyanates. 

Isomeric  compounds.  It  will  be  noted  that  ammonium  cyanate  and 
urea  have  the  same  molecular  formula,  namely,  N2H4CO.  Compounds 
like  these,  which  have  the  same  molecular  formula,  are  known  as 
isomeric  compounds,  or  simply  as  isomers.  Their  difference  in  proper- 
ties is  due  to  the  different  arrangement  of  the  atoms  in  the  molecule. 
A  great  many  isomeric  compounds  are  known,  especially  among  the 
compounds  of  carbon. 

THE  HYDROCARBONS  AND  SOME  OF   THEIR  SIMPLE 
DERIVATIVES 

Carbon  and  hydrogen  combine  to  form  a  large  number  of  compounds 
known  collectively  as  the  hydrocarbons.  For  convenience  these  com- 
pounds are  divided  into  a  number  of  groups,  or  series,  each  one  being 
named  from  its  first  member.  In  the  table  below  are  given  the  names 
and  formulas  of  a  few  of  the  simpler  members  of  the  four  most  im- 
portant groups.  It  will  be  noted  that  the  members  in  each  group  are 
arranged  in  accordance  with  the  number  of  carbon  atoms  present.  The 
general  formula  for  the  members  of  each  group  is  added,  in  which  the 
letter  n  represents  the  number  of  carbon  atoms.  The  methane  group 
is  the  most  extensive,  all  the  compounds  up  to  C^H^  being  known. 

METHANE  SERIES  ETHYLENE  SERIES  BENZENE  SERIES 

CH4       ...     methane  C2H4   .     .     .     ethylene         C6H6      .     .     .     benzene 

C2H6     .     .     .     ethane  C3H6   .     .     .     propylene      C7H8       .     .     .     toluene 

C3H8     .     .     .     propane  C4H8   .     .     .     butylene        C8H10     .     .     .     xylene 

C4H10    .     ,     .     butane  CnH2n  CnH2n_6 

C6H12    .     .     .     pentane 

C6H14    .     .     .     hexane  ACETYLENE  SERIES 

CnH2n  +  2  C2H2   .     .     .     acetylene 
C3H4  .     .     .     allylene 

Homologous  series.  It  will  be  noticed  that  the  formulas  of  the  successive 
members  of  each  of  the  above  series  differ  by  the  group  of  atoms  CH2.  Such  a 
series  is  called  a  homologous  series.  In  general  it  may  be  stated  that  the  mem- 
bers of  a  homologous  series  show  a  regular  gradation  in  most  physical  properties 
and  are  similar  in  chemical  properties.  The  boiling  points,  for  example,  gradually 


CARBON  AND  ITS  COMPOUNDS  293 

increase  with  the  number  of  carbon  atoms  present.  Thus,  in  the  methane  series  the 
boiling  points  are  such  that  under  ordinary  conditions  of  temperature  and  pressure 
the  first  four  members  are  gases ;  those  containing  from  five  to  sixteen  carbon 
atoms  are  liquids,  the  boiling  points  of  which  increase  with  the  number  of  carbon 
atoms  present ;  those  containing  more  than  sixteen  carbon  atoms  are  solids. 

Structural  formulas  of  the  hydrocarbons.  The  structural  formula 
of  the  first  member  of  each  of  the  above  series  is  as  follows : 

CH 

H  /^ 

1  TT  IT  HC         CH 

H-C-H  £>C  =  C<Jt  H-C=C-H  II        I 

HC    CH 

H  \// 

CH 

Methane  (CH4)      Ethylene  (C2H4)       Acetylene  (C2H2)       Benzene  (C6H6) 

The  hydrocarbons  belonging  to  the  methane  series  differ  from  those 
of  the  other  series  in  that  they  do  not  unite  directly  with  any  other 
element  or  radical.  In  the  case  of  methane  this  property  is  indicated 
by  the  structural  formula  assigned  above,  which  represents  the  carbon 
atom  as  combined  with  four  different  hydrogen  atoms.  Since  carbon 
is  never  known  to  have  a  valence  greater  than  4,  it  is  evident,  in 
accordance  with  the  formula,  that  methane  will  not  combine  directly 
with  any  other  element,  for  to  do  so  would  be  to  increase  the  valence 
of  carbon  above  this  value.  Such  compounds  are  known  as  saturated 
compounds.  While  other  elements  do  not  combine  directly  with 
saturated  compounds,  they  may  be  substituted  for  elements  already 
present.  Thus,  when  chlorine  acts  upon  methane  under  suitable  con- 
ditions, one  of  the  hydrogen  atoms  is  displaced  by  a  chlorine  atom, 
forming  a  compound  of  the  formula  CHgCl.  Compounds  like  ethy- 
lene,  on  the  other  hand,  are  called  umaturated,  for  they  combine 
directly  with  certain  elements.  Thus,  ethylene  combines  directly  with 
chlorine,  forming  a  compound  of  the  formula  C2H4C12.  It  might  seem 
that  this  property  of  ethylene  could  best  be  expressed  by  the  formula 

TT  TT 

H>C— C<H»  in  which  each  carbon  atom  is  trivalent.  Since  carbon 
tends  to  act  as  a  tetravalent  element,  one  would  expect  a  compound 
of  this  character  to  add  directly  the  two  atoms  of  chlorine,  each  car- 
bon atom  becoming  thereby  tetravalent,  as  expressed  in  the  formula 
Clx  /Cl 

H-C— C-H.  It  seems  probable,  however,  that  the  carbon  atoms  in 
K'  XH 

ethylene  are  really  tetravalent,  and  this  fact  is  expressed  by  the 
double  union,  or,  as  it  is  often  termed,  the  double  bond,  between  the  two 


294  GENEEAL  CHEMISTRY 

carbon  atoms.    The  addition  of  chlorine  is  then  expressed  as  follows : 

H  C1\          /C1 

£J>C  =  C<;;  +  CL     =     H— C— C— H.   Similarly,  in  acetylene  there  is  a 

H7          ^U 

triple  bond  between  the  two  carbon  atoms,  and  such  a  compound  may 
add  either  two  or  four  atoms  of  a  univalent  element. 

Sources  of  the  hydrocarbons.   There  are  two  chief  sources  of  the 
hydrocarbons,  namely,  petroleum  and  coal  tar. 

1.  Petroleum.  This  liquid  is  pumped  from  wells  driven  into  the 
earth  in  certain  localities.  California,  Oklahoma,  Texas,  and  Pennsyl- 
vania are  the  chief  oil-producing  regions  in  the  United  States.  The 
crude  petroleum  consists  largely  of  liquid  hydrocarbons,  in  which  are 
dissolved  both  gaseous  and  solid  hydrocarbons.    For  most  purposes 
it  is  refined  before  it  is  used.    In  this  process  the  petroleum  is  run 
into  large   iron  stills  and  subjected  to  fractional  distillation.    The 
various  hydrocarbons  distill  over  in  the  general  order  of  their  boiling 
points.    The  distillates  which  collect  between  certain  limits  of  tem- 
perature  are  kept  separate   and   serve  for  different  uses ;  they   are 
further  purified,  generally  by  washing  first  with  su  If  uric  acid,  then 
with  an  alkali,  and  finally  with  water. 

Among  the  products  obtained  in  the  distillation  of  petroleum  are 
the  following,  named  in  the  general  order  of  their  boiling  points : 
pentane,  hexane,  the  naphthas,  kerosene,  or  coal  oil,  lubricating  oils, 
vaseline,  and  paraffin.  Pentane  is  used  as  the  standard  illuminant  in 
determining  the  candle  power  of  flames.  Hexane  is  used  as  a  solvent. 
A  number  of  different  naphthas  are  recognized  commercially,  differing 
in  boiling  point  and  density.  The  naphthas  of  low  boiling  point  are 
used  as  fuels  in  gasoline  stoves  and  in  motors  ;  those  of  higher  boiling 
point  are  used  in  the  manufacture  of  paints.  The  terms  benzine  and 
.gasoline  are  often  applied  in  a  general  way  to  the  most  common  of 
the  naphthas.  It  must  be  remembered  that,  with  the  exception  of 
pentane  and  hexane,  none  of  these  products  are  definite  chemical  com- 
pounds. Each  consists  of  a  mixture  of  hydrocarbons,  the  boiling  points 
of  which  lie  within  certain  limits. 

2.  Coal  tar.  This  product  is  obtained  in  the  manufacture  of  coal 
gas  (p.  323).   It  is  a  complex  mixture  and  is  refined  by  the  same  gen- 
eral method  used  in  refining  petroleum.    The  principal  hydrocarbons 
obtained  from  the  coal  tar  are  benzene,  toluene,  naphthalene,  and 
anthracene.    In  addition  to  the  hydrocarbons,  coal  tar  contains  many 
other  compounds,  such  as  carbolic  acid  (phenol)  and  aniline. 


CARBON  AND  ITS  COMPOUNDS  295 

Properties  of  the  hydrocarbons.  The  hydrocarbons  are  all  readily  in- 
flammable, the  carbon  and  hydrogen  present  combining  with  oxygen 
to  form  carbon  dioxide  and  water  respectively.  The  members  of  the 
methane  series  are  very  stable  compounds  and,  with  the  exception  of 
oxygen,  even  the  most  active  reagents  have  little  effect  upon  them. 
The  members  of  the  other  series  are  much  more  reactive.  It  is  advis- 
able here  to  discuss  only  a  very  few  of  the  individual  hydrocarbons. 

Methane  (marsh  gas)  (CHJ.  This  hydrocarbon  constitutes  about 
90  per  cent  of  natural  gas.  It  is  formed  in  marshes  by  the  decay  of 
vegetable  matter  under  water,  and  bubbles  of  the  gas  are  often  seen 
to  rise  when  the  dead  leaves  on  the  bottom  of  pools  are  stirred.  It 
also  collects  in  mines,  and,  when  mixed  with  air,  is  called  fire  damp 
by  the  miners,  because  of  its  great  inflammability,  damp  being  an  old 
name  for  a  gas.  It  is  formed  when  organic  matter,  such  as  coal  or 
wood,  is  heated  in  closed  vessels,  and  is  therefore  a  principal  con- 
stituent of  coal  gas. 

Methane  is  prepared  in  the  laboratory  by  heating  sodium  acetate 
with  soda  lime.  The  latter  substance  is  a  mixture  of  sodium  and 
calcium  hydroxides.  Regarding  it  as  sodium  hydroxide  alone,  the 
equation  for  the  reaction  is  as  follows : 

NaC2H302  +  NaOH  =  Na2CO3  +  CH4 

Methane  is  a  colorless,  odorless  gas  0.55  times  as  heavy  as  air.  It 
is  but  very  slightly  soluble  in  water.  It  can  be  condensed  to  a  color- 
less liquid  which  boils  at  —  164°  under  a  pressure  of  1  atmosphere. 
It  burns  with  a  pale-blue  flame,  its  heat  of  combustion  amounting  to 
211,930  cal. 

Ethane  (C2H6).  This  hydrocarbon  can  be  prepared  from  methane  by  first  dis- 
placing an  atom  of  hydrogen  by  one  of  chlorine,  and  then  treating  the  resulting 
compound  with  sodium : 

2  CH3C1  +  2  Na  =  2  NaCl  +  C2H6 

Ethane  is  similar  to  methane  in  properties. 

Ethylene  (C^H^).  Small  amounts  of  ethylene  are  present  in  coal  gas.  It  is 
prepared  by  the  action  of  sulf  uric  or  phosphoric  acid  on  alcohol.  In  the  reaction 
the  alcohol  loses  the  elements  of  water : 

C2H60  =  C2H4  +  H20 

Acetylene  (C2H2).  This  is  a  colorless  gas  and  is  formed  by  the  direct 
combination  of  carbon  and  hydrogen  at  very  high  temperatures.  It  is 
also  formed  when  certain  hydrocarbons  are  burned  in  a  limited  supply 


296  GENERAL  CHEMISTRY 

of  air,  so  that  the  combustion  is  incomplete.  In  this  way  it  is  formed 
when  the  flame  of  a  Bunsen  burner  "  strikes  back,"  that  is,  when  the 
flame  bums  at  the  bottom  of  the  tube.  The  easiest  as  well  as  the 
most  economical  method  for  its  preparation  consists  in  the  action  of 
water  upon  calcium  carbide : 

CaC2  +  H20:=CaO+C2H2 

Pure  acetylene  is  an  odorless  gas.  As  ordinarily  prepared,  however, 
it  has  a  disagreeable  odor  due  to  impurities.  It  is  0.92  times  as  heavy 
as  air.  At  a  temperature  of  0°  it  is  condensed  to  a  colorless  liquid  by 
a  pressure  of  26.05  atmospheres.  At  lower  temperatures  it  forms  a 
solid  melting  at  —  81°.  Under  ordinary  conditions  it  burns  with  a  very 
smoky  flame  due  to  the  incomplete  supply  of  oxygen.  In  burners  so 
constructed  as  to  secure  a  large  admixture  of  air  it  burns  with  a  bril- 
liant white  light.  It  is  an  endothermic  compound.  In  the  formation 
of  a  gram-molecular  weight  of  the  gas  48,200  cal.  are  absorbed.  When 
the  compound  is  decomposed,  this  same  quantity  of  heat  is  evolved. 
When  acetylene  is  burned,  this  heat  of  decomposition  is  added  to  the 
heat  generated  by  the  combustion  of  the  carbon  and  hydrogen,  so 
that  the  total  quantity  of  heat  evolved  is  very 
great,  amounting  to  310,000  cal.  for  each  gram- 
molecular  weight  of  the  gas  burned.  While 
acetylene  is  stable  at  high  temperatures,  it  is 
very  explosive  at  ordinary  temperatures  when 
under  pressure,  and  many  accidents  have  resulted 
,A  from  attempts  to  condense  the  gas  in  cylinders 

for  commercial  use. 

The  preparation  and  combustion  of  acetylene.  The  gas 
can  be  prepared  in  a  generator  such  as  is  shown  in  Fig.  104. 
The  inner  tube  A  contains  lumps  of  calcium  carbide, 

while  the  outer  one  is  filled  with  water.   As  long  as  the 

FIG.  104  stopcock  B  is  closed,  the  water  cannot  rise  in  the  inner 

tube.   When  the  stopcock  is  open,  the  water  rises  and, 

coming  into  contact  with  the  carbide,  generates  acetylene.  The  gas  escapes 
through  the  stopcock.  After  the  air  has  been  expelled  from  the  interior  of  the 
tube,  the  gas  may  be  lighted  as  it  issues  from  the  burner. 

Uses  of  acetylene.  Acetylene  is  used  both  as  an  illuminant  and  as 
a  source  of  intense  heat.  As  an  illuminant  it  is  used  especially  in 
isolated  places,  where  neither  gas  nor  electric  lights  are  available.  It 
has  been  found  that  the  gas  can  be  compressed  with  safety  by  forcing 


CABBON  AND  ITS  COMPOUNDS          297 

it  at  low  temperatures  into  metal  cylinders  completely  filled  with  some 
porous  material  (such  as  a  mixture  of  asbestos  and  cotton),  which  is 
partially  saturated  with  acetone  or  acetaldehyde.  These  liquids  absorb 
large  volumes  of  the  gas,  and  under  the  conditions  it  is  nonexplosive. 
Stored  in  this  way  the  gas  is  now  a  common  article  of  commerce. 

The  intense  heat  generated  by  the  combustion  of  acetylene  makes 
it  useful  in  certain  processes  requiring  high  temperatures,  such  as  the 
welding  and  cutting  of  metals.  For  this  purpose  the  acetylene  is 
burned  in  an  apparatus  known  as  the  oxyacetylene  blowpipe,  which 
is  exactly  like  the  oxyhydrogen  blowpipe.  A  temperature  of  about 
2700°  may  be  obtained  in  this  way.  This  blowpipe  has  been  found 
especially  useful  in  dismantling  iron  structures,  such  as  the  battleship 
Maine,  since  the  tip  of  the  flame,  when  drawn  slowly  over  the  metal, 
burns  it  at  the  point  of  contact  and  thus  makes  it  possible  to  cut  the 
metal  into  pieces. 

Benzene.  This  hydrocarbon  is  obtained  commercially  from  coal  tar. 
It  is  a  colorless  liquid  boiling  at  80.2°,  and  is  a  good  solvent  for  most 
organic  compounds,  so  that  it  is  very  useful  in  the  laboratory.  Both 
the  liquid  and  its  vapor  are  highly  inflammable.  It  differs  from  the 
other  hydrocarbons  studied  in  that  it  readily  reacts  with  nitric  acid. 
The  product  of  the  reaction  is  nitrobenzene  (CgHgNOg),  a  slightly 
yellowish  liquid  often  called  oil  of  mirbane: 


When  nitrobenzene  is  reduced  with  hydrogen,  a  nearly  colorless  liquid 
known  as  aniline  is  formed.    This  has  the  formula  O  H.NH    and  is 

65  2 

the  compound  from  which  some  of  the  aniline  dyes  are  prepared. 

Naphthalene  and  anthracene.  These  hydrocarbons  occur,  along  with  benzene,  in 
coal  tar.  They  are  solids,  insoluble  in  water.  The  well-known  moth  balls  are 
made  of  naphthalene.  Large  quantities  of  naphthalene  are  used  in  the  prepara- 
tion of  indigo,  a  dye  formerly  obtained  entirely  from  the  indigo  plant  grown  in 
India,  but  now  prepared  by  laboratory  methods.  Similarly,  anthracene  is  used 
in  the  preparation  of  the  dye  alizarin.  This  dye  was  formerly  obtained  from 
the  root  of  the  madder  plant,  which  was  extensively  cultivated  for  this  purpose, 
especially  in  France. 

Substitution  products  of  the  hydrocarbons.  As  a  rule,  at  least  a  part 
of  the  hydrogen  in  any  hydrocarbon  can  be  displaced  by  certain  ele- 
ments or  groups  of  elements.  Thus,  the  compounds  CH3C1,  CH2C12, 
CHC13,  and  CC14  can  be  obtained  from  methane  by  the  action  of 


298  GENERAL  CHEMISTRY 

chlorine.  Such  compounds  are  called  substitution  products.  Among 
the  important  substitution  products  of  methane  are  the  following : 

1.  Chloroform  (CHC13).   This  is  the  well-known  compound  used  as  an 
anesthetic  in  surgery.    It  is  a  colorless,  heavy  liquid  boiling  at  61°. 

2.  lodoform  (CHI3).  It  is  a  yellow,  crystalline  solid,  largely  used  as 
an  antiseptic. 

3.  Carbon  tetrachloride  (CC/4).    This  compound  is  a  heavy,  colorless, 
oily  liquid  boiling  at  76.7°.    Like  chloroform  and  benzene,  it  is  a  good 
solvent  for  many  organic  compounds,  such  as  the  fats,  and  is  used  for 
this  purpose,  as,  for  example,  in  removing  grease  spots  from  fabrics. 

CARBOHYDRATES 

The  term  carbohydrate  is  applied  to  a  class  of  compounds  which 
includes  the  sugars,  starch,  and  allied  bodies.  These  compounds  con- 
tain carbon,  hydrogen,  and  oxygen ;  the  last  two  elements  are  usually 
present  in  the  ratio  in  which  they  combine  to  form  water.  The  most 
important  members  of  this  class  are  the  following : 

Sucrose  (cane  sugar) C12H22On 

Lactose  (milk  sugar) Ci2H22°ii '  H2° 

Maltose .  C12H22On  -  H2O 

Dextrose  (grape  sugar) C6H12O6 

Levulose  (fruit  sugar) C6H12O6 

Starch (C6H1005)* 

Cellulose    ...........  (C0H10O6)X 

The  molecular  formulas  of  starch  and  cellulose  are  unknown,  but  they 
are  known  to  be  multiples  of  the  formula  C6H10Og,  and  hence  are  rep- 
resented in  the  above  table  by  (C6H10O5).,..  In  the  discussion  of  these 
compounds  they  will  be  represented  by  the  simple  formula  C6HloOg. 

Sucrose  (C12H22On).  This  is  the  substance  ordinarily  called  sugar. 
It  occurs  in  many  plants,  especially  in  the  sugar  cane  and  sugar  beet, 
each  of  which,  at  the  present  time,  furnishes  approximately  50  per 
cent  of  the  total  production. 

When  a  solution  of  sucrose  is  heated  to  about  70°  with  hydro- 
chloric acid,  two  isomeric  sugars,  dextrose  and  levulose,  are  formed 
in  accordance  with  tjie  following  equation: 

C12H22On  +  H20  =  C6H]206  +  C6HJ2Oe 

When  heated  to  160°,  sucrose  melts ;  if  the  temperature  is  increased 
to  about  215°,  a  partial  decomposition  takes  place  and  a  brown  sub- 
stance, known  as  caramel,  forms.  This  is  used  as  a  coloring  matter. 


CARBON  AND  ITS  COMPOUNDS  299 

Lactose  (milk  sugar)  (C12H22011  •  H20).  This  sugar  is  isomeric  with 
sucrose  and  is  present  in  the  milk  of  all  mammals.  The  average  com- 
position of  cow's  milk  is  as  follows  : 

Water    .............  87.17% 

Casein  (nitrogenous  matter)       .....  3.56% 

Butter  fat  ............  3.64% 

Lactose       ............  4.88% 

Mineral  matter    ..........  0.75% 

When  rennin,  a  substance  obtained  from  the  stomach  of  calves,  is 
added  to  milk,  the  casein  separates  and  is  used  in  the  manufacture  of 
cheese.  The  remaining  liquid  contains  the  lactose,  which  separates  on 
evaporation.  Lactose  resembles  sucrose  in  appearance,  but  is  not  so 
sweet  or  soluble.  The  souring  of  milk  is  due  to  the  fact  that  the 
lactose  present  is  changed  into  lactic  acid.  The  acid  gives  to  the 
milk  its  sour  taste,  and  also  causes  the  separation  of  the  casein,  thus 
producing  the  well-known  appearance  of  sour  milk. 

Maltose  (C12H22011  •  H20).  This  sugar  resembles  sucrose  and  lactose 
in  its  general  properties.  It  is  prepared  by  the  action  of  malt  upon 
starch  ;  hence  the  name  maltose.  Malt  is  the  name  applied  to  barley 
which  has  been  moistened,  kept  in  a  warm  place  until  it  has  germinated, 
and  then  heated  until  the  vitality  of  the  grain  has  been  destroyed.  In 
the  process  of  germination  a  substance  is  formed  known  as  diastase,  and 
it  is  this  substance  which  imparts  to  malt  its  property  of  changing  starch 
into  maltose.  It  is  from  this  sugar  that  alcohol  and  most  of  the  alcoholic 
liquors  are  prepared. 

Dextrose  (grape  sugar)  (C6H1206).  Dextrose  is  present  in  many  fruits 
and  is  commonly  called  grape  sugar  because  of  its  presence  in  grape 
juice.  It  can  be  obtained,  along  with  levulose,  by  heating  sucrose  with 
acids.  Commercially  it  is  prepared  by  heating  starch  with  dilute 
hydrochloric  acid.  The  acid  acts  simply  as  a  catalytic  agent,  the  re- 
action really  taking  place  between  the  starch  and  water.  The  starch 
is  first  changed  into  a  sweet-tasting  solid  known  as  dextrin,  and  this, 
on  further  action,  is  transformed  into  dextrose  : 


Pure  dextrose  is  a  white,  crystalline  solid  readily  soluble  in  water, 
and  is  not  so  sweet  as  sucrose.  It  is  prepared  in  large  quantities 
and,  being  less  expensive,  is  used  as  a  substitute  for  sucrose  in  the 
manufacture  of  jellies,  jams,  molasses,  candy,  and  other  sweets.  As 


300  GENERAL  CHEMISTRY 

sold  on  the  market  it  is  usually  in  the  form  of  a  thick,  colorless  sirup 
known  as  glucose,  or  corn  sirup.  This  sirup  contains  from  40  to  50  per 
cent  dextrose,  from  30  to  40  per  cent  dextrin,  and  water. 

Levulose  (fruit  sugar)  (C6H1206).  Levulose  is  a  white  solid  which 
occurs  along  with  dextrose  in  fruits  and  honey.  It  is  sweet  and  has 
the  general  properties  of  a  sugar. 

Cellulose  (C6H1005).  Cellulose  forms  the  basis  of  all  woody  fibers. 
Cotton  and  linen  are  nearly  pure  cellulose.  It  is  insoluble  in  water, 
alcohol,  and  dilute  acids.  Nitric  acid  reacts  with  it,  forming  explosive 
nitrates  variously  known  as  nitrocellulose,  pyroxylin,  and  guncotton. 
When  exploded,  they  yield  only  colorless  gases ;  hence  they  are  used 
in  the  manufacture  of  smokeless  gunpowder.  Collodion  is  a  solution 
of  nitrocellulose  in  a  mixture  of  alcohol  and  ether.  Celluloid  is  a  mix- 
ture of  nitrocellulose  and  camphor.  Paper  consists  mainly  of  cellulose, 
the  finer  grades  being  made  from  linen  and  cotton  rags,  and  the  cheaper 
grades  from  straw  and  wood. 

Starch  (C6H100B).  This  compound  is  by  far  the  most  abundant  car- 
bohydrate found  in  nature,  being  present  especially  in  seeds  .and 
tubers.  In  the  United  States  it  is  obtained  chiefly  from  corn,  .over  70 
per  cent  of  which  is  starch.  In  Europe  it  is  obtained  principally  from 
the  potato.  Starch  consists  of  minute  granules.  These  granules  are 
composed  of  a  substance  known  as  granulose,  surrounded  by  a  mem- 
brane composed  principally  of  cellulose.  The  granulose  is  soluble 

in  water.  Starch  does 
not  dissolve  in  cold 
water,  however,  since 
the  granulose  is  pro- 
tected from  the  action 
of  water  by  the  in- 
soluble cellulose  mem- 
brane. When  heated 
with  water,  the  mem- 
Fio.  105.  Corn  starch  FIG.  106.  Wheat  starch  branes  burst  and  the 
Magnified  260  diameters  Magnified  260  diameters  ,  T  i 

granulose    dissolves. 

Starch  granules  differ  somewhat  in  appearance,  according  to  the  source 
of  the  starch,  so  that  it  is  generally  possible  to  determine  the  origin 
of  any  particular  sample  by  its  microscopic  appearance.  Figs.  105  and 
106  represent  the  appearance  of  typical  granules  of  starch  derived 
from  corn  and  wheat,  when  viewed  under  the  microscope. 


CAKBON  AND  ITS  COMPOUNDS  301 

ALCOHOLS 

The  alcohols  may  be  regarded  as  derived  from  the  hydrocarbons  by 
substituting  for  one  or  more  hydrogen  atoms  a  corresponding  number 
of  hydroxyl  groups.  A  great  many  alcohols  are  known,  and,  like  the 
hydrocarbons,  may  be  arranged  in  series.  The  relation  between  the 
first  three  members  of  the  methane  series  and  the  corresponding 
alcohols  is  shown  in  the  following  table: 

CH4    (methane)    ....     CH3OH    (methyl  alcohol) 
C2H6  (ethane)  .     ....     C2H5OH  (ethyl  alcohol) 
C3H8  (propane)     ....     C3H7OH  (propyl  alcohol) 

The  terms  methyl,  ethyl,  and  propyl,  used  in  designating  the  differ- 
ent alcohols,  are  names  applied  to  the  univalent  radicals  CH3,  C2H5, 
and  CgH7  respectively.  It  will  be  noted  that  the  names  of  these  radi- 
cals are  derived  from  the  names  of  the  corresponding  hydrocarbons 
by  changing  the  ending  -ane  to  -yl. 

Methyl  alcohol  (wood  alcohol)  (CH3OH).  This  compound  is  obtained 
in  the  destructive  distillation  of  wood,  and  on  this  account  is  called 
wood  alcohol.  It  is  a  colorless  liquid  which  has  a  density  of  0.79 
and  boils  at  64.7°.  It  burns  with  an  almost  colorless  flame  and  is 
sometimes  used  for  heating  purposes  in  place  of  the  more  expensive 
ethyl  alcohol.  It  is  a  good  solvent  for  organic  substances  and  is  used 
to  a  considerable  extent  as  a  solvent  in  the  manufacture  of  varnishes. 
It  is  quite  poisonous.  It  has  a  specific  action  upon  the  optic  nerve, 
and  many  cases  of  blindness  have  resulted  both  from  drinking  the 
liquid  and  from  repeatedly  inhaling  its  vapor. 

Duncan  states  that  "  out  of  ten  men  who  drink  4  oz.  of  pure  methyl  alcohol  in 
any  form  whatever,  four  will  probably  die,  two  of  them  becoming  blind  before 
death.  The  remaining  six  may  recover,  but  of  these,  two  will  probably  be 
permanently  blind." 

Ethyl  alcohol  (C2HBOH).  This  is  the  compound  commonly  designated 
as  alcohol. 

1.  Preparation.  It  is  prepared  by  the  action  of  ordinary  brewers' 
yeast  upon  certain  sugars,  especially  maltose  and  dextrose.  With 
dextrose  the  reaction  is  expressed  by  the  following  equation : 

C.HU0.=  2C,H.OH+.2CO, 

This  process,  in  which  a  sugar  is  changed  into  alcohol  and  carbon 
dioxide  by  the  action  of  yeast,  is  known  as  alcoholic  fermentation.    The 


302 


GENERAL  CHEMISTRY 


yeast  is  a  low  form  of  plant  life  which  grows  in  the  sugar  solution 
under  suitable  conditions.  During  its  growth  it  secretes  a  substance 
known  as  zymase,  which  is  the  active  agent  in  effecting  alcoholic  fer- 
mentation. While  sucrose  does  not  ferment  directly,  the  addition  of 
yeast  to  its  aqueous  solution  first  resolves  the  sucrose  into  dextrose 
and  levulose,  both  of  which  then  ferment. 

Laboratory  preparation  of  alcohol.  The  formation  of  alcohol  and  carbon  dioxide 
from  dextrose  may  be  shown  as  follows :  About  100  g.  of  dextrose  is  dissolved 

in  a  liter  of  water  in  the 
flask  A  (Fig.  107).  This 
flask  is  connected  with  the 
bottle  B,  which  is  partially 
filled  with  limewater.  The 
tube  C  contains  solid  so- 
dium hydroxide.  A  little 
bakers'  yeast  is  now  added 
to  the  solution  in  flask  A, 
and  the  apparatus  is  con- 
nected as  shown  in  the 
figure.  If  the  temperature 
FIG.  107  is  maintained  at  about  30°, 

the   reaction    soon   begins. 

The  bubbles  of  gas  escape  through  the  limewater  in  B.  A  precipitate  of  calcium 
carbonate  soon  forms  in  the  limewater,  showing  the  presence  of  carbon  dioxide. 
The  sodium  hydroxide  in  tube  C  prevents  the  carbon  dioxide  in  the  air  from 
acting  upon  the  limewater.  The  alcohol  remains  in  the  flask  A  and  may  be 
separated  by  fractional  distillation. 

Commercially,  alcohol  is  prepared  from  starch  obtained  chiefly  from 
corn  and  potatoes.  The  starch  is  first  converted  into  maltose  by  the 
action  of  malt,  the  maltose  being  then  changed  into  alcohol  and  carbon 
dioxide  by  yeast.  In  this  way  it  is  possible  to  obtain  an  aqueous 
solution  containing  from  15  to  20  per  cent  of  alcohol.  By  fractional 
distillation  this  may  be  concentrated  to  a  solution  containing  96  per 
cent  of  alcohol.  When  lime  (CaO)  is  added  to  this  solution,  and  the 
mixture  heated,  most  of  the  remaining  water  combines  with  the  lime 
to  form  calcium  hydroxide.  Upon  distilling  the  resulting  mixture, 
alcohol  containing  less  than  1  per  cent  of  water  distills  over.  Such 
alcohol  is  termed  absolute  alcohol.  The  ordinary  alcohol  of  the  drug- 
gist contains  approximately  95  per  cent  by  volume  of  alcohol. 

2.  Properties.  Ethyl  alcohol  is  a  colorless  liquid  with  a  pleasant 
odor.  It  has  a  density  of  0.789  at  20°,  boils  at  78.30°,  and  solidi- 
fies at  —  112.3°.  It  resembles  methyl  alcohol  in  its  general  properties 


CAKBON  AND  ITS  COMPOUNDS  303 

and  is  sometimes  used  as  a  source  of  heat,  since  its  flame  is  very  hot 
and  does  not  deposit  carbon,  as  does  the  flame  from  oils.  When  taken 
into  the  system  in  small  quantities  it  causes  intoxication;  in  large 
quantities  it  acts  as  a  poison.  The  intoxicating  properties  of  such 
liquors  as  beer,  wine,  and  whisky  are  due  to  the  alcohol  present. 
When  heated  to  140°  with  sulfuric  acid  it  loses  the  elements  of  water, 
forming  ordinary  ether,  as  shown  in  the  following  equation : 

C,H5OfH]     C,H       Q 
C2H5[0  H|     C2H5^ 

Ether  is  largely  used  as  an  anesthetic  in  surgical  operations. 

Denatured  alcohol.  The  federal  government  imposes  a  heavy  tax  on  alcohol 
and  alcoholic  liquors,  the  exact  amount  of  the  tax  varying  according  to  the  per- 
centage of  alcohol  present.  For  the  95  per  cent  alcohol  this  tax  is  $2.11  per  gal- 
lon. This  increases  the  cost  of  the  alcohol  to  such  an  extent  that  it  is  not 
economical  to  use  it  for  many  purposes  for  which  it  is  adapted,  such  as  for  a 
solvent  in  the  preparation  of  paints  and  varnishes  and  as  a  material  for  the 
preparation  of  many  important  organic  compounds.  By  an  act  of  Congress  in 
1906  the  tax  was  removed  from  denatured  alcohol,  that  is,  from  alcohol  mixed 
with  some  substance  which  renders  it  unfit  for  the  preparation  of  a  beverage  but 
which  does  not  impair  its  value  for  manufacturing  purposes.  Some  of  the  Euro- 
pean countries  have  similar  laws.  The  substances  ordinarily  used  as  denaturants 
are  wood  alcohol,  gasoline,  and  pyridine. 

Alcoholic  liquors.  All  alcoholic  liquors  are  made  by  alcoholic  fermentation. 
Wine  is  made  by  the  fermentation  of  the  dextrose  in  grape  juice  and  contains 
from  5  to  15  per  cent  by  volume  of  alcohol.  Beer  is  made  from  maltose  formed 
by  the  action  of  malt  upon  starch  obtained  from  various  grains,  chiefly  barley. 
It  contains  from  3  to  5  per  cent  by  volume  of  alcohol.  Whisky  contains  about 
50  per  cent  by  volume  of  alcohol  and  is  made  from  starch  by  a  process  very 
similar  to  that  described  under  the  commercial  preparation  of  alcohol.  Almost 
any  saccharine  liquid,  such  as  cider  and  the  juices  of  fruits  in  general,  gradually 
undergoes  alcoholic  fermentation  when  exposed  to  air,  the  yeast  cells  entering 
from  the  air. 

Glycerin  (C3H8(OH)3).  This  compound  may  be  regarded  as  derived 
from  propane  (C3Hg)  by  replacing  three  atoms  of  hydrogen  with  three 
hydroxyl  groups.  It  is  therefore  an  alcohol.  It  is  an  oily,  colorless 
liquid  having  a  sweetish  taste,  and  is  obtained  in  the  manufacture 
of  soaps.  Glycerin  is  used  in  medicine  and  in  the  manufacture  of 
nitroglycerin  and  dynamite. 

Nitroglycerin  and  dynamite.  Xitric  acid  reacts  with  glycerin  in  the  same  way 
that  it  reacts  with  any  base  containing  three  hydroxyl  groups,  such  as  Fe(OH)3: 

Fe  (OH)3  +  3  HN03  -  Fe  (NO3)3  +  3  H2O 
C3H5(OH)3  +  3  HN08  =  C3Hfi(N03)3  +  3  H2O 


304  GENERAL  CHEMISTRY 

The  resulting  nitrate,  C3Hg  (NO3)3,  is  the  main  constituent  of  nitroglycerin,  a 
slightly  yellowish  oil  characterized  by  its  explosive  properties.  It  explodes  by 
pressure,  by  detonation,  or  by  heating  to  250°.  The  following  equation  represents 
in  a  general  way  the  changes  which  take  place  in  the  decomposition  of  nitro- 
glycerin :  4  c^  ^NO^g  =  12  CQ2  +  6  NS  +  10  H2Q  +  Q2 

One  volume  of  nitroglycerin  on  explosion  yields  about  1300  volumes  of  gaseous 
compounds,  which  are  expanded  by  the  heat  of  explosion  to  over  10,000  volumes. 
Dynamite  consists  of  a  mixture  of  sodium  nitrate,  wood  pulp,  and  nitroglycerin. 
The  wood  pulp  acts  as  an  absorbent  for  the  nitroglycerin.  The  strength  of  the 
dynamite  depends  on  the  percentage  of  nitroglycerin  present.  Dynamite  is 
used  much  more  than  nitroglycerin,  since  it  does  not  explode  so  readily  on  per- 
cussion, and  therefore  can  be  transported  with  safety. 

ALDEHYDES 

When  treated  with  suitable  oxidizing  agents,  alcohols  are  converted, 
by  loss  of  hydrogen,  into  compounds  known  as  aldehydes.  The  most 
important  aldehyde  is  formaldehyde  (CH2O),  which  is  prepared  by  the 
oxidation  of  methyl  alcohol  : 

2  CH3OH  +  O2  =  2  CH2O  +  2  H2O 

Formaldehyde  is  a  gas  and  is  largely  used  as  a  disinfectant.  An 
aqueous  solution  containing  40  per  cent  by  weight  of  the  gas  is  sold 
by  druggists  under  the  name  formalin.  When  oxidized,  formalde- 
hyde yields  formic  acid  (CH2O2)  : 


ACIDS 

Like  the  other  classes  of  organic  compounds,  the  organic  acids  may 
be  arranged  in  homologous  series.  One  of  the  most  important  of  these 
series  is  the  fatty-add  series,  so  called  because  the  derivatives  of  certain 
of  its  members  are  constituents  of  the  fats.  Some  of  the  most  impor- 
tant members  of  the  series  are  given  in  the  following  table.  They  are 
all  monobasic,  and  this  fact  is  expressed  in  the  formulas  by  separating 
the  replaceable  hydrogen  atom  from  the  rest  of  the  molecule  : 


o 

o 


H  •  CHO2     ....  formic  acid,  a  liquid  boiling  at  100° 

H  •  C2H3O2  ....  acetic  acid,  a  liquid  boiling  at  118' 

H  •  C3H5O2  ....  propionic  acid,  a  liquid  boiling  at  140 

H  •  C4H7O2  ....  butyric  acid,  a  liquid  boiling  at  163 

H  •  C16H31O2     .     .     .  palmitic  acid,  a  solid  melting  at  62° 

H  •  C18H35O2     .     .     .  stearic  acid,  a  solid  melting  at  69° 


CARBON  AND  ITS  COMPOUNDS          305 

Formic  acid  (H'CH02).  This  is  a  colorless  liquid  and  occurs  in 
many  plants,  such  as  the  stinging  nettle.  It  is  also  present  in  a  certain 
species  of  ant ;  hence  the  name  formic  acid,  the  word  formic  being 
derived  from  a  Latin  word  meaning  "  ant." 

Acetic  acid  (H*C2H302).  This  is  best  known  as  the  acid  which  im- 
parts the  sour  taste  to  vinegar.  It  is  prepared  commercially  by  the 
destructive  distillation  of  wood  (p.  280).  It  is  a  colorless  liquid 
and  has  a  strong,  pungent  odor.  When  anhydrous,  it  crystallizes  as 
a  white  solid  which  melts  at  18°,  and  closely  resembles  ice  in  appear- 
ance ;  hence  the  name  glacial  acetic  acid.  Many  of  the  salts  of  acetic 
acid  are  well-known  compounds.  Thus,  lead  acetate  (Pb  (C2H3O2)2) 
is  the  white  solid  known  as  sugar  of  lead. 

Vinegar.  All  vinegars  are  prepared  by  the  action  of  the  vegetable  organism, 
known  as  Mycoderma  aceti,  or,  commonly,  as  "  mother  of  vinegar,"  upon  a  liquid 
containing  alcohol.  This  change  of  alcohol  into  acetic  acid  through  the  action 
of  the  organism  is  known  as  acetic  fermentation,  and  may  be  expressed  by  the 
following  equation  :  ^^  +  ^  =  R  ^^  +  ^ 

Instead  of  starting  with  the  alcoholic  liquid  one  may  use  some  substance  which 
contains  starch  or  sugar,  such  as  barley,  cider,  or  molasses.  In  such  cases  the 
starch  and  sugar  present  are  first  converted  into  alcohol,  as  explained  under  the 
preparation  of  alcohol.  Thus,  in  the  manufacture  of  cider  vinegar  the  sugar 
in  the  apple  juice  first  undergoes  alcoholic  fermentation,  the  alcohol  formed 
then  undergoing  acetic  fermentation.  Vinegars  contain  from  4  to  6  per  cent  of 
acetic  acid. 

Butyric  acid  (H»C4H702).  ^Butyric  acid  is  a  liquid  of  disagreeable 
odor.  A  derivative  of  the  acid  is  present  in  butter  and  gives  it  its 
characteristic  taste. 

Palmitic  and  stearic  acids.  These  are  white  solids  insoluble  in  water. 
They  are  obtained  from  fats  (p.  306). 

Relation  between  the  hydrocarbons,  alcohols,  aldehydes,  and  acids.  The  state- 
ment has  been  made  that  the  alcohols  may  be  regarded  as  derived  from  the  hydro- 
carbons by  the  substitution  of  a  hydroxyl  group  for  hydrogen ;  also,  that  the 
alcohols,  when  oxidized,  yield  aldehydes  by  loss  of  hydrogen,  and  that  the  alde- 
hydes on  further  oxidation  are  changed  into  acids.  The  relation  between  these 
compounds  may  be  seen  from  the  following  formulas,  representing  methane  and 
ethane  and  the  corresponding  alcohols,  aldehydes,  and  acids. 


CH4 

(methane) 

CH3OH 

(methyl  alcohol) 

CH20 

(formaldehyde) 

CH2O2 
(formic  acid) 

C2H6 

(ethane) 

C2H6OH 

(ethyl  alcohol) 

C2H4O 
(acetaldehyde) 

C2H402 

(acetic  acid) 

306  GENERAL  CHEMISTRY 

Acids  belonging  to  other  series.  In  addition  to  the  members  of  the 
series  described  above,  mention  may  be  made  of  the  following  well- 
known  acids,  the  first  four  of  which  are  white  solids : 

1.  Oxalic  acid  (H2C20^-2H20)  is  found  in  many  plants. 

2.  Malic  acid  (H^CJI^O^-H^O)  occurs  in  a  free  state  in  apples,  pears, 
and  other  fruits,  as  well  as  in  the  berries  of  the  mountain  ash. 

3.  Tartaric  acid  (H2-C^06)  occurs  in  many  fruits,  especially  the 
grape,  either  in  a  free  state  or  in  the  form  of  its  salts.    Some  of 
its  salts  are  well-known   compounds.    The  potassium  acid  tartrate 
(KHC4H4O6)  is  ordinarily  known  as  cream  of  tartar.    It  is  a  white 
solid  obtained  from  grape  juice  in  the  manufacture  of  wine,  and  is 
used  in  the  manufacture  of  baking  powders.    Potassium  sodium  tar- 
trate   (KNaC4H4O6  •  4  H2O)   is    used   in    medicine    under   the    name 
Rochelle  salt. 

4.  Citric  acid  (H3-C6H607'H20)  occurs  especially  in  lemons. 

5.  Lactic  acid  (H-C3H503)  is  a  liquid  formed  from  lactose  in  the 
souring  of  milk.    The  formation  of  lactic  acid  from  lactose  is  known 
as  lactic  fermentation ;   like  alcoholic  and  acetic  fermentation,  it  is 
caused  by  the  presence  of  a  low  form  of  vegetable  organism. 

6.  Oleic  acid  (H-  C18fl"3302)  is  an  oily  liquid.    Certain  derivatives  of 
oleic  acid  constitute  the  principal  part  of  many  oils  and  fats. 

ESTERS,  OILS,  AND  FATS 

When  acids  are  mixed  with  alcohols  under  certain  conditions,  a 
reaction  occurs  similar  to  that  which  takes  place  between  acids  and 
bases.  The  following  equations  will  serve  as  illustrations : 

KOH  +  HN03  =  KN03  +  H2O 

CH3OH  +  HNO3  +=±.  CH3NO3  +  H2O 

CH3OH  +  H  •  C2H302  q=±  CH3  -  C2H302  +  H20 

The  resulting  compounds,  of  which  methyl  acetate  (CH3C2H3O2)  may 
be  taken  as  an  example,  are  known  as  esters.  They  differ  from  ordi- 
nary salts  in  that  they  contain  a  hydrocarbon  radical,  such  as  CHg  or 
C2H5,  in  place  of  a  metal. 

The  fats  are  largely  mixtures  of  the  esters  known  as  olein,  palmitin, 
and  stearin.  These  esters  may  be  regarded  as  derived  from  oleic,  pal- 
mitic, and  stearic  acids  respectively  by  replacing  the  hydrogen  of  the 
acid  with  the  glycerin  radical  C8H6.  This  radical  is  trivalent,  and  since 
oleic,  palmitic,  and  stearic  acids  contain  only  one  replaceable  hydrogen 


CARBON  AND  ITS  COMPOUNDS  307 

atom  to  the  molecule,  it  is  evident  that  three  molecules  of  each  acid 
must  enter  into  the  formation  of  each  molecule  of  the  ester.  The 
formulas  for  the  acids  and  the  esters  derived  from  each  are  as  follows : 


ACIDS 

ESTEBS 

H  •  C18H83O2  (oleic  acid) 
H  •  C16H31O2  (palmitic  acid) 
H  •  C18H35O2  (stearic  acid) 

C3H5(C18H3302)3(olein) 
C3H5(C16H3102)3  (palmitin) 
C3H5(C18H3502)3  (stearin) 

Olein  is  a  liquid,  while  palmitin  and  stearin  are  solids.  The  oils  are 
mainly  olein,  while  the  solid  fats  are  mainly  palmitin  and  stearin. 

Butter  fat  and  oleomargarine.  Butter  fat  consists  principally  of  olein,  palmitin, 
and  stearin.  The  flavor  of  the  fat  is  due  to  the  presence  of  a  small  percentage  of 
butyrin,  which  is  an  ester  of  butyric  acid  and  has  the  composition  C3H5(C4H7O2)3. 
Oleomargarine  differs  from  butter  mainly  in  that  a  smaller  amount  of  butyrin  is 
present.  It  is  made  from  fat  obtained  from  cattle  and  hogs.  Small  percentages 
of  cottonseed  oil  are  also  sometimes  used.  This  fat  is  churned  with  milk,  or  a 
small  quantity  of  butter  is  added,  in  order  to  furnish  sufficient  butyrin  to  impart 
the  butter  flavor. 

Saponification ;  soaps.  When  an  ester  such  as  ethyl  nitrate  (C2H&NO3) 
is  heated  with  an  alkali,  a  reaction  expressed  by  the  following  equation 
takes  place :  C2H5NO3  +  KOH  =  C2H5OH  +  KNO3 

This  type  of  reaction  is  known  as  saponification,  since  it  is  the  one 
which  takes  place  in  the  manufacture  of  soaps.  The  ordinary  soaps 
are  made  by  heating  fats  with  a  solution  of  sodium  hydroxide.  The 
reactions  involved  may  be  illustrated  by  the  following  equation  rep- 
resenting the  reaction  between  palmitin  and  sodium  hydroxide : 

C,H6(CleH3102)3  +  3  NaOH  =  3  NaCMHnO,  +  C8H5(OH)3 

In  accordance  with  this  equation  the  esters  which  constitute  the  fats 
and  oils  are  converted  into  glycerin  and  the  sodium  salts  of  the  corre- 
sponding acids.  The  sodium  salts  are  separated  and  constitute  ordinary 
soaps.  These  salts  are  soluble  in  water.  When  added  to  water  con- 
taining calcium  salts,  the  insoluble  calcium  palmitate  and  stearate  are 
precipitated.  Magnesium  salts  act  in  a  similar  way.  It  is  because  of 
these  facts  that  so  much  soap  is  used  up  by  hard  water.  The  glycerin 
formed  in  the  manufacture  of  soaps  is  recovered,  and  it  is  from  this 
source  that  the  glycerin  of  commerce  is  obtained. 

The  proteins.  The  term  protein  is  applied  to  a  large  class  of  com- 
plex nitrogenous  compounds  which  are  everywhere  abundant  in  all 


308 


GENERAL   CHEMISTRY 


animal  and  vegetable  organisms  and  constitute  the  principal  part  of 
the  tissues  of  the  living  cell.  They  all  contain  nitrogen,  carbon,  hydro- 
gen, and  oxygen,  and  some  also  contain  sulfur  and  phosphorus. 

Foods.  While  the  compounds  present  in  our  foods  are  very  numer- 
ous and  often  exceedingly  complex,  yet  they  may  all  be  included  in 
a  few  general  classes.  It  is  customary  to  regard  the  edible  portion  of 
our  foods  as  composed  of  proteins,  fats,  carbohydrates,  mineral  matter, 
and  water.  Since  the  mineral  matter  is  left  as  a  residue  when  the  food 
is  burned,  it  is  listed  as  ash,  in  reporting  the  analyses  of  foods. 

In  a  general  way  it  may  be  stated  that  the  protein  matter  in  our 
food  serves  to  replace  the  worn-out  tissues  of  our  bodies,  as  well  as  to 
supply  material  for  growth.  The  carbohydrates  and  fats  are  more 
or  less  interchangeable,  since  they  are  both  oxidized  in  the  body 
(p.  121),  and  thus  serve  as  a  source  of  heat  and  muscular  energy. 
The  mineral  matter  supplies  the  material  for  building  up  the  solid 
tissues  of  the  body  and  has,  in  addition,  other  more  complex  functions. 
The  protein  matter  may  fulfill  the  same  function  as  the  fats  and  carbo- 
hydrates if  the  latter  are  lacking  in  our  foods.  Since  the  various  con- 
stituents of  our  foods  serve  different  purposes,  it  is  evident  that  a 
proper  mixture  of  these  is  essential  to  health. 

The  composition  of  the  edible  portion  of  a  few  typical  foods  is 
given  in  the  following  table,  taken  from  Sherman's  "  Chemistry  of 
Food  and  Nutrition." 


WATER 

(Per  cent) 

PROTEIN 

(Per  cent) 

FAT 

(Per  cent) 

CARBO- 
HYDRATES 
(Per  cent) 

ASH 

(Per  cent) 

Beef  free  from  visible  fat 
Ham,  smoked,  lean 
Salmon 

73.8 
53.5 
64.6 

22.1 
20.2 
212 

2.9 
20.8 
12  8 

1.2 

5.5 
14 

Eggs  
Milk  . 

73.7 
87.0 

14.8 
3  3 

10.5 
4  0 

5  0 

1.0 

0  7 

Butter  

li.o 

1.0 

85.0 

30 

Oatmeal  .  .  .'  . 

7.3 

16.1 

72 

67.5 

1  9 

Rice  

12.3 

80 

0  3 

79  0 

04 

Wheat,  flour  .... 
Bread,  white  .... 
Beans,  dried  .... 
Corn,  green  .  .  . 
Potatoes  .  .  .  . 

11.9 
35.3 
12.6 
75.4 

78.3 

13.3 
9.2 
22.5 
3.1 
2  2 

1.5 
1.3 

1.8 
1.1 
0  1 

72.7 
53.1 
59.6 
19.7 
18  4 

0.6 
1.1 
3.5 
0.7 
1  0 

Tomatoes  . 

94.3 

0  9 

04 

3  9 

0  5 

Apples  . 

84.6 

04 

0  5 

14  2 

0  3 

CHAPTER  XXII 

MOLECULAR  WEIGHTS 

Introduction.  It  was  shown  in  Chapter  VII  that  from  the  results  of 
the  analysis  of  a  compound  it  is  possible  to  calculate  the  simplest 
formula  which  correctly  represents  its  composition.  It  was  assumed, 
however,  that  we  already  have  a  concordant  system  of  atomic  weights. 
Moreover,  it  was  pointed  out  that  the  formula  so  calculated  is  merely 
the  simplest  one  possible,  and  does  not  necessarily  represent  the 
composition  of  the  molecule.  Thus,  the  simplest  formula  for  hydro- 
gen peroxide  is  HO,  but  there  is  good  reason  for  concluding  that' 
the  molecular  formula  is  really  H2O2.  The  composition  of  hydro- 
fluoric acid  is  satisfactorily  expressed  by  the  formula  HF,  but  the 
fact  that  it  forms  the  acid  salt  KHF2  suggests  the  double  formula 
H2F2  as  more  correctly  representing  the  molecule.  It  is  the  purpose 
of  the  present  chapter  to  develop  the  methods  by  which  the  true 
molecular  formulas  of  compounds  may  be  determined,  and  by  which 
a  concordant  system  of  atomic  weights  may  be  deduced. 

Methods  for  securing  equal  numbers  of'  molecules  of  different  com- 
pounds. If  it  were  possible  to  actually  count  out  equal  numbers  of 
molecules  of  various  compounds  into  separate  piles,  it  is  evident  that 
the  ratio  between  the  weights  of  the  several  piles  would  be  the  same 
as  the  ratio  between  the  weights  of  the  individual  molecules.  If  one 
of  the  piles  were  to  be  taken  as  standard,  it  would  then  be  possible  to 
state  how  much  heavier  each  kind  of  molecule  is  than  those  of  the 
standard  pile,  and  such  figures  would  be  the  relative  weights  of  the 
various  molecules. 

Evidently  this  cannot  be  done  directly,  but  it  has  been  found  pos- 
sible to  accomplish  the  same  results  indirectly  through  the  discovery 
that  certain  properties  of  substances  are  dependent  merely  upon  the 
number  of  molecules  present,  quite  irrespective  of  their  character. 
This  provides  a  ready  means  for  deciding  when  we  are  dealing  with 
the  same  number  of  molecules  of  different  substances,  and  conse- 
quently for  determining  their  relative  weights.  Several  of  these 
properties  will  now  be  discussed. 

309 


310  GENERAL  CHEMISTEY 

1.  The  volume  of  a  gas  is  proportional  to  the  number  of  molecules  which 
it  contains.    There  are  many  reasons  for  believing  that  this  statement 
is  true,  though  it  is  evident  that  its  truth  cannot  be  experimentally 
demonstrated.    Historically  its  acceptance  was  based  upon  the  formu- 
lation of  a  very  important  law,  together  with  an  hypothesis  as  to  the 
meaning  of  the  law. 

The  law  of  Gay-Lussac.  In  the  early  years  of  the  nineteenth  century 
the  Frenchman  Gay-Lussac  investigated  the  proportion  by  volume  in 
which  gases  combine,  as  well  as  the  relation  between  their  individual 
volumes  and  that  of  the  product  formed.  His  studies  brought  to  light 
the  remarkable  relationships  illustrated  in  the  following  equations : 

1  volume  hydrogen  +  1  volume  chlorine  =  2  volumes  hydrogen  chloride 

2  volumes  hydrogen  +  1  volume  oxygen  =  2  volumes  steam 

3  volumes  hydrogen  +  1  volume  nitrogen  =  2  volumes  ammonia 

2  volumes  carbon  monoxide  +  1  volume  oxygen  =  2  volumes  carbon  dioxide 

The  conclusions  which  he  reached,  verified  by  all  subsequent  research, 
may  be  stated  in  the  following  form,  known  as  the  law  of  Gay-Lussac : 
When  two  gases  combine,  there  is  an  integer  ratio  between  their  vol- 
umes, as  well  as  between  the  volume  of  either  of  them  and  that  of  the 
product,  provided  it  is  a  gas. 

Avogadro's  hypothesis.  In  1811  Avogadro,  professor  of  physics  at 
Turin,  suggested  that  the  most  probable  explanation  of  this  striking 
generalization  is  that  equal  volumes  of  all  gases  (under  the  same 
physical  conditions)  contain  the  same  number  of  molecules.  This 
suggestion  is  still  called  an  hypothesis,  because  it  cannot  be  directly 
verified  by  experiment.  It  is  in  complete  accord  with  the  various  gas 
laws,  and  can  be  shown  to  be  a  logical  conclusion  if  we  accept  the 
kinetic  theory  of  gases. 

Assuming  the  truth  of  this  hypothesis,  we  have  at  once  a  means  of 
setting  apart  an  equal  number  of  molecules  of  various  gases,  for  they 
will  be  contained  in  any  definite  volume  which  we  may  choose,  for 
example  in  1 1.  The  ratios  between  the  weights  of  1  1.  of  each  of  the 
various  gases  (measured  under  the  same  conditions)  will  be  the  same 
as  those  between  the  weights  of  the  several  kinds  of  molecules. 

2.  The  lowering  of  the  freezing  point  of  a  solvent  is  proportional  to  the 
molecular  concentration  of  the  solute.    In  the  chapter  on  solutions  it  was 
shown  that  the  lowering  of  the  freezing  point  of  a  solution  depends 
not  upon  the  kind  of  molecules  dissolved  in  the  solvent,  but  merely 
upon  their  concentration  in  it,  provided  there  is  no  ionization  (law 


MOLECULAR  WEIGHTS  811 

of  Raoult).  Conversely,  when  the  freezing  point  of  a  given  quantity 
of  a  solvent  is  lowered  to  the  same  extent  by  two  different  substances, 
we  have  an  equal  number  of  molecules  of  solute  present  in  the  two 
cases.  If,  therefore,  a  definite  quantity  of  a  solvent  is  taken,  say  1 1., 
and  experiments  are  made  to  determine  the  quantities  of  various  sul> 
stances  which  must  be  dissolved  in  it  to  produce  a  definite  lowering 
of  the  freezing  point,  say  of  1°,  then  these  quantities  will  contain  an 
equal  number  of  molecules,  and  the  ratio  between  their  weights  will 
be  the  same  as  that  between  the  weights  of  the  individual  molecules. 

3.  Change  in  boiling  point  and  vapor  pressure.  In  an  entirely  similar  way  the  boil- 
ing point  and  the  vapor  pressure  of  liquids  have  been  found  to  be  changed  to  the 
same  extent  by  the  same  number  of  dissolved  molecules,  irrespective  of  their 
character,  and  by  measurement  of  these  changes  it  is  possible  to  determine  the 
relative  weights  of  the  molecules  producing  them. 

The  standard  for  molecular  weights.  Having  devised  methods  for 
the  determination  of  the  relative  weights  of  molecules,  the  next  step 
is  to  agree  upon  some  one  substance  as  a  standard,  so  as  to  express 
these  weights  in  multiples  of  that  of  the  standard  molecule.  Since 
molecular  weights  are  most  frequently  determined  by  measurements 
of  the  volumes  of  gases,  it  is  best  to  select  as  a  standard  some 
gaseous  substance,  preferably  an  element.  Various  gases  have  been 
chosen  at  different  times,  the  choice  being  guided  merely  by  con- 
venience. Hydrogen  commends  itself  as  being  the  lightest  of  all 
gaseous  substances,  so  that,  if  it  is  taken  as  standard,  all  others 
will  have  weights  greater  than  unity.  On  the  whole,  oxygen  serves 
as  the  most  satisfactory  standard,  and  is  now  universally  adopted. 
We  need  only  determine  the  weight  of  a  liter  of  each  kind  of  gas, 
and  state  how  much  heavier  it  is  than  that  of  a  liter  of  oxygen,  to 
have  a  series  of  molecular  weights  based  on  oxygen  as  unity. 

Weight  of  oxygen  taken  as  32.  The  assumption  of  oxygen  as  unity 
is  open  to  the  objection  that  a  number  of  gases  are  lighter  than  oxygen, 
and  then-  molecular  weights  would,  on  this  basis,  be  less  than  unity, 
which  would  be  undesirable.  Hydrogen  is  the  lightest  of  all,  the  ratio 
between  equal  volumes  of  hydrogen  and  oxygen  being  1 :  15.87.  The 
smallest  whole  number  assignable  to  oxygen  which  will  at  the  same 
time  place  hydrogen  as  great  as  unity  is  therefore  16. 

In  adopting  any  standard  for  molecular  weights,  however,  it  should 
be  remembered  that  the  same  standard  must  serve  for  both  mole- 
cules and  atoms.  If  there  should  be  any  reason  for  thinking  that 


312  GEKEKAL  CHEMISTEY 

the  molecule  of  hydrogen  consists  of  more  than  one  atom,  it  would  be 
better  to  adopt  a  still  higher  value  for  the  oxygen  molecule,  so  as 
to  place  the  hydrogen  atom  at  a  value  as  great  as  unity.  That  the 
hydrogen  molecule  consists  of  at  least  two  atoms  is  readily  shown 
by  the  following  reasoning.  When  hydrogen  combines  with  chlorine, 
the  volume  relations  are  expressed  by  the  equation 

1  volume  H  + 1  volume  Cl  =  2  volumes  HC1 

According  to  Avogadro's  hypothesis  the  two  volumes  of  hydrogen 
chloride  must  contain  twice  as  many  molecules  as  the  one  volume  of 
either  hydrogen  or  chlorine.  But  each  of  the  molecules  of  hydrogen 
chloride  must  contain  at  least  one  atom  of  hydrogen.  This  accounts 
for  twice  as  many  atoms  of  hydrogen  in  the  hydrogen  chloride  as 
there  are  molecules  in  the  one  volume  of  hydrogen.  These  relations 
can  be  brought  into  harmony  with  the  hypothesis  of  Avogadro  by 
assuming  that  each  molecule  of  hydrogen  is  made  up  of  two  atoms. 

It  will  be  noticed  that  this  reasoning  merely  shows  that  there  are 
twice  as  many  hydrogen  atoms  in  the  hydrogen  molecules  as  there 
are  in  those  of  hydrogen  chloride.  There  might  be  two  in  each  of  the 
latter  and  four  in  the  former,  but  since  there  are  no  facts  known 
which  point  to  the  larger  numbers,  we  make  the  simplest  assumption 
possible,  and  conclude  that  the  numbers  are  two  and  one. 

Since  it  is  the  atom  of  hydrogen,  rather  than  its  molecule,  which 
we  wish  to  hold  as  great  as  unity,  and  since  it  appears  that  there  are 
two  atoms  in  the  molecule,  we  shall  have  to  double  the  value  16,  which 
we  have  provisionally  adopted  for  the  oxygen  molecule,  and  place  it 
at  32.  This  will  give  the  value  2.016  to  the  hydrogen  molecule  and 
1.008  to  the  hydrogen  atom. 

The  gram-molecular  volume.  Avogadro's  hypothesis  states  that 
equal  volumes  of  gases  contain  the  same  number  of  molecules.  If, 
then,  the  volume  of  oxygen  which  contains  32  g.  of  the  gas  is  deter- 
mined, it  is  clear  that  the  number  stating  the  weight  in  grams  of  any 
other  gas  which  occupies  this  same  volume  will  also  tell  how  much 
heavier  the  gas  is  than  oxygen  taken  as  32.  In  other  words,  it  will  be 
the  molecular  weight  of  the  gas  referred  to  oxygen  as  32.  Accurate 
experiment  has  shown  that  1  1.  of  oxygen  weighs  1.429  g.,  so  that 
32  g.  will  occupy  32  -f- 1.429  =  22.38,  or,  in  round  numbers,  22.4  1. 
If  we  construct  a  vessel  of  exactly  this  capacity,  it  will  hold  an 
equal  number  of  molecules  of  any  other  gas,  so  that  the  weight  of 


MOLECULAR  WEIGHTS 


313 


this  volume  of  the  gas  will  express  its  molecular  weight  compared 
with  that  of  oxygen  taken  as  32.  This  volume  is  therefore  called  the 
gram-molecular  volume.  We  thus  reach  the  general  rule :  To  determine 
the  molecular  weight  of  a  gas  referred  to  oxygen  as  32,  find  the 
weight  in  grams  of  22.4  1.  of  the  gas.  This  process  is  termed  the 
determination  of  the  vapor  density  of  the  substance,  referred  to 
oxygen  as  32. 

Experimental  determinations  of  molecular  weights.  In  an  actual  ex- 
periment we  determine  the  exact  weight  of  any  convenient  volume 
of  a  gas  under  any  convenient  conditions  of  temperature  and  pres- 
sure, and  from  this  weight  calculate  the  weight  of  22.4  1.  under 
standard  conditions.  Two  general  methods  for  making  such  a  deter- 
mination are  in  use. 

1.  Method  of  Dumas.    This  older  method  for  determining  vapor  densities,  em- 
ployed by  Dumas  as  early  as  1827,  is  readily  adapted  to  the  present  purpose.   A 
small  glass  bulb  of  about  100  cc.  capacity  is  attached  to  an  air  pump,  exhausted 
of  the  air  which  it  contains,  and  weighed  empty ;  or  its  volume  is  determined, 
and  from  its  weight  in  air  the  weight  of  the  air 

which  it  contains  is  subtracted,  to  give  its  real 
weight  when  empty.  The  flask  is  then  filled  with 
the  gas  under  investigation,  and  again  weighed. 
If  we  know  the  volume  of  the  flask,  the  weight 
of  the  gas  which  fills  it,  and  the  temperature  and 
pressure  under  which  it  was  filled,  it  is  easy  to 
calculate  the  weight  of  22.4  1.  of  the  gas  under 
standard  conditions.  The  method  is  also  adapted 
to  easily  volatile  liquids.  In  such  cases  the  liquid 
whose  molecular  weight  is  to  be  determined  may 
be  placed  in  the  flask  and  the  flask  immersed  in 
a  bath,  the  temperature  of  which  is  above  the 
boiling  point  of  the  liquid  (Fig.  108).  The  liquid 
rapidly  boils  away,  leaving  the  flask  filled  with 
the  vapor  of  the  liquid  at  the  temperature  of  the 
bath  and  the  pressure,  of  the  atmosphere.  The 
flask  may  then  be  closed  and  its  weight  determined.  From  the  results  of  this 
experiment,  together  with  the  volume  and  weight  of  the  empty  flask,  the  molecu- 
lar weight  of  the  liquid  in  the  vapor  state  may  be  deduced. 

2.  Method  of  Victor  Meyer.    In  1878  Victor  Meyer  devised  a  more,  convenient 
though  less  accurate  method,  which  depends  upon  a  somewhat  different  principle. 
The  apparatus  employed  is  represented  in  Fig.  109.    The  inner  vessel  A  is  a  long, 
narrow  tube  expanded  into  a  bulb  at  the  lower  end,  open  at  the  upper  end  B, 
and  furnished  with  a  slender  side  tube  C  near  the  top.    This  vessel  is  placed  in 
an  outer  jacket  Z),  which  contains  some  liquid,  frequently  water,  by  boiling  which 
the  inner  vessel  may  be  raised  to  a  definite,  steady  temperature.    When  this  has 


FIG.  108 


314 


GENERAL   CHEMISTRY 


i 


been  effected,  a  graduated  collecting  tube  E  is  filled  with  water  and  inverted 
over  the  end  of  the  delivery  tube  C,  as  shown  in  the  figure.  A  small  quantity 
of  a  liquid  (from  0.1  to  0.2  g.)  whose  molecular  weight  is  to  be  determined  is 
weighed  out  in  a  minute  bottle  F,  and  the  bottle  is  dropped 
in  at  B,  the  opening  being  quickly  closed  by  a  stopper.  The 
liquid  in  the  jacket  must  have  a  higher  boiling  point  than 
that  whose  molecular  weight  is  to  be  determined.  The 
liquid  in  the  bottle  will  then  be  very  rapidly  converted 
into  vapor  in  the  bulb,  and  an  equal  volume  of  air  from 
the  upper  part  of  the  vessel  A  will  be  forced  over  into  the 
measuring  tube  E,  which  can  then  be  measured  at  room 
temperature  and  calculated  to  standard  conditions.  Since 
all  gases  undergo  the  same  volume  changes  when  tempera- 
ture and  pressure  are  altered,  it  is  evident  that  this  calcu- 
lation gives  us  the  volume  which  the  vapor  of  the  weighed 
substance  would  occupy  under  standard  conditions  if  it 
were  possible  to  so  obtain  it.  From  the  weight  and  the 
volume  it  is  easy  to  calculate  the  weight  of  22.4  1. 

Freezing-point  method.  There  are,  however,  many 
substances  whose  vapor  density  cannot  be  meas- 
ured. In  some  cases  the  temperature  of  volatiliza- 
tion is  so  high  that  containing  vessels  which  will 
stand  such  a  temperature  cannot  be  constructed. 
In  other  cases  the  substance  decomposes  before 
vaporizing,  as  is  true  with  most  metallic  salts  and 
a  great  many  organic  substances,  such  as  sugar. 
In  many  of  these  cases  it  is  possible  to  determine 
the  molecular  weight  by  measurements  of  the  lowering  of  the  freezing 
point  of  some  suitable  solvent.  The  first  step  is  to  determine  the  low- 
ering produced  by  a  gram-molecular  weight  of  some  substance  whose 
molecular  weight  is  known  from  the  use  of  vapor-density  methods. 
For  example,  such  methods  give  us  the  formula  C2H6O  for  alcohol, 
with  a  molecular  weight  of  46.  If  46  g.  of  alcohol  is  dissolved  in 
a  liter  of  water,  it  is  found  that  the  freezing  point  of  the  water  is 
lowered  by  1.87°.  According  to  the  law  of  Raoult  (p.  133)  a  gram- 
molecular  weight  of  any  substance  dissolved  in  a  liter  of  water  will 
lower  the  freezing  point  by  the  same  amount,  provided  only  that  it  is 
not  ionized  or  changed  chemically.  It  will  be  seen  that  in  a  general 
way  the  determination  of  this  value  of  molecular  lowering  (1.87°)  cor- 
responds to  the  fixing  of  the  gram-molecular  volume  for  gases  (22.4  1.). 
To  determine  the  molecular  weight  of  some  substance  of  unknown 
value  we  need  only  dissolve  a  weighed  quantity  in  a  definite  volume 


FXG.  109 


MOLECULAR  WEIGHTS 


315 


of  water  and  determine  the  lowering  of  the  freezing  point.  The  cal- 
culation of  the  molecular  weight  is  then  very  simple.  Thus,  0.46  g. 
of  sugar,  dissolved  in  20.35  g.  of  water,  produced  a  lowering  of  0.126°. 
The  proportion  20.35  : 1000  : :  0.46  :  x 

x  =  22.604 

gives  the  weight  which  would  produce  the  same  lowering  if  dissolved 
in  a  liter  of  water.  But  a  gram-molecular  quantity  would  produce  a 

lowering  of  1.87°,  so  that   22.604  g.  constitutes  -ir-^r  of  a  gram- 

1.87 

molecular  weight  of  sugar.  The  molecular  weight  as  determined  by 
this  experiment  is  therefore  335.4  (the  correct  weight  being  342). 

The  apparatus  usually  employed  in  such  measure- 
ments, known  as  the  Beckmann  apparatus,  is  repre- 
sented in  Fig.  110.  The  weighed  solvent  (water)  is 
placed  in  the  inner  tube  A,  and  the  thermometer 
arranged  to  dip  into  the  liquid.  A  suitable  stirrer  is 
also  provided.  The  tube  A  is  hung  in  a  larger  tube 
B,  which  is  empty,  and  the  latter  is  surrounded  by 
a  cooling  mixture  of  ice  and  salt.  This  cools  the  air 
in  B  below  the  freezing  point  of  water,  and  makes 
it  possible  to  freeze  a  portion  of  the  solution  in  A. 
The  thermometer  is  of  special  construction,  so  that 
changes  of  temperature  as  'small  as  0.001°  can  be 
read.  By  noting  the  freezing  point  of  the  pure  water, 
and  also  the  freezing  point  after  the  introduction  of 
a  weighed  quantity  of  some  substance,  the  data  for 
the  calculation  are  readily  obtained. 

Calculation  of  formulas  from  molecular 
weights.  Having  devised  satisfactory  meth- 
ods for  determining  molecular  weights,  it  is 
a  simple  matter  to  calculate  the  correct  molec- 
ular formula.  Thus,  if  analysis  shows  that 
the  composition  of  sulfur  chloride  is  correctly 
represented  by  the  formula  SCI,  having  a  formula  weight  of  67*53, 
while  measurement  of  its  vapor  density  gives  the  value  133.2  as  its 
approximate  molecular  weight,  it  is  evident  that  the  simple  formula 
must  be  doubled  to  obtain  the  molecular  formula.  This  gives  the 
formula  S2C12,  with  a  molecular  weight  of  135.06. 

In  general  the  molecular  formula  is  found  by  the  following  pro- 
cedure: (1)  determine  the  simplest  formula  by  analysis ;  (2)  deter- 
mine the  approximate  molecular  weight;  (3)  multiply  the  simplest 


FIG.  110 


316  GENEKAL  CHEMISTRY 

formula  by  the  integer  which  will  give  a  value  near  to  the  experi- 
mentally determined  molecular  weight. 

Atomic  weights.  It  will  be  recalled  that  the  mere  analysis  of  com- 
pounds enables  us  to  determine  the  combining  weights  of  the  elements, 
but  does  not  tell  which  multiple  of  the  simplest  combining  weight 
really  represents  the  relative  weight  of  the  atom.  Having  developed 
satisfactory  methods  for  determining  the  molecular  weights  of  com- 
pounds, it  is  now  an. easy  matter  to  , decide  definitely  what  multiple 
of  the  combining  weight  should  be  chosen  as  the  atomic  weight. 

Approximate  atomic  weights  deduced  from  molecular  weights.  The 
method  by  which  this  is  done  can  best  be  explained  by  some  examples. 
In  the  table  on  the  opposite  page  the  first  column  gives  the  names  of  a 
number  of  compounds  containing  oxygen,  hydrogen,  carbon,  nitrogen, 
chlorine,  and  sulfur.  The  second  column  gives  the  molecular  weights 
of  the  compounds  as  determined  by  one  of  the  methods  just  described. 
These  values  are  merely  approximate,  being  subject,  as  a  rule,  to  an 
error  of  several  per  cent.  The  succeeding  columns  show  how  many 
of  the  units  of  the  molecular  weight  must  be  assigned  to  the  several 
atoms  composing  the  molecules.  These  values  are  obtained  by  care- 
fully analyzing  the  compounds  and  then  multiplying  the  molecular 
weights  by  the  percentage  of  each  element  present.  Thus,  if  the  molecu- 
lar weight  of  carbon  monoxide  is  approximately  27,  and  analysis  shows 
that  the  substance  contains  42.96  per  cent  of  carbon,  it  is  clear  that, 
of  the  27  units  constituting  the  molecular  weight,  27  X  0.4296  =  11.6 
must  be  assigned  to  the  carbon  atoms  present. 

Now,  each  molecule  must  be  made  up  of  some  definite  number  of 
each  kind  of  atoms  composing  it,  so  that  in  the  molecular  weight  of 
the  compound  the  part  by  weight  assigned  to  each  atom  must  repre- 
sent either  the  relative  weight  of  the  atom  or  some  multiple  of  it.  In 
a  considerable  number  of  compounds  of  a  given  element  the  molecules 
of  some  of  these  will,  in  all  probability,  contain  but  a  single  one  of 
these  atoms.  In  such  compounds  the  part  of  the  molecular  weight 
assigned  to  the  atom  in  question  will  be  approximately  its  real  atomic 
weight.  In  all  other  compounds  containing  this  element  the  parts 
assigned  to  its  atom  will  be  some  multiple  of  this  smallest  weight. 
An  examination  of  the  several  columns  in  the  table  will  show  that 
this  is  the  case.  In  each  column  all  the  values  listed  are  approximate 
multiples  of  the  smallest  one,  which  is  indicated  in  round  numbers  at 
the  bottom  of  the  column.  These  values  are  approximate  only,  for  the 


MOLECULAR  WEIGHTS 


317 


reason  that  the  molecular  weights  are  not  accurate,  and  so  the  values 
derived  from  them  are  subject  to  the  same  error. 

For  the  determination  of  the  approximate  atomic  weight  of  an 
element  we  therefore  reach  the  following  procedure:  (1)  determine 
the  molecular  weight  of  a  large  number  of  compounds  of  the  element 
in  question ;  (2)  analyze  these  compounds ;  (3)  multiply  the  molec- 
ular weight  of  each  by  the  percentage  of  the  element  present  in  the 
compound.  The  least  value  so  obtained  will  be  the  approximate 
atomic  weight. 

TABLE  ILLUSTRATING  CALCULATION  OF  ATOMIC  WEIGHTS 


NAME 

MOL. 
WEIGHT 

PART 
0 

PART 
H 

PART 
C 

PART 
N 

PART 

Cl 

PART 

s 

FORMULA 

Carbon  monoxide 

27.. 

154 

11.6 

CO 

Carbon  dioxide      .     . 

44.4 

32.2 

12.2 

C02. 

Methane        .... 

16.5 

4.1 

12.4 

CH4 

Acetylene     .... 

25.8 

2.2 

23.6 

C2H2 

Benzene  .     .     .     .  '  . 

77.6 

5.8 

71.8 

C  H  * 

Alcohol    

46.6 

16.1 

6.3 

24.2 

cXo 

Susrar 

340.0 

174.7 

22.2 

143.1 

C    H    O 

Hydrogen  chloride     . 

36.7 

1.1 

35.6 

HC1  ^ 

Carbon  tetrachloride 

153.0 

11.9 

142.1 

CC1, 

Chloroform  .... 

121.0 

1.1 

12.1 

107.7 

CHC13 

Sulfur  chloride 

133.2 

70.1 

63.1 

S2C12 

Sulfuryl  chloride   . 

136.2 

32.4 

71.5 

32.3 

S02C12 

Sulfur  dioxide 

64.15 

32.85 

31.3 

S02 

Carbon  disulfide    .     . 

'  76.9 

12.3 

64.7 

CS2 

Ammonia     .... 

17.3 

3.1 

14.2 

XH8 

Nitrous  oxide   .     .     . 

43.5 

15.9 

27.6 

N2O 

Nitrogen  pentoxide    . 

107.1 

79.3 

27.8 

^2^5 

Nitric  acid  . 

63.75 

48.4 

1.05 

14.3 

HNOS 

Water      .     .     . 

17.9 

15.9 

2.0 

H20 

Hydrogen  peroxide    . 

34.3 

32.2 

2.1 

H2O2 

Approximate  atomic  weight 

16.00 

1.0 

12.0 

14.0 

35.5 

32.0 

The  accurate  determination  of  atomic  weights.  For  exact  deter- 
minations of  atomic  weights  we  must  now  return  to  the  combining 
weights  or  equivalents  (p.  .89).  These  can  be  determined  with  great 
precision  by  the  analysis  of  suitable  compounds,  and  they  are  always 
either  identical  with  the  atomic  weights  or  bear  a  simple  integer  rela- 
tion to  them.  Thus,  from  the  above  table  it  will  be  seen  that  the 
approximate  atomic  weight  of  sulfur  is  32.  The  equivalent  of  sulfur, 
as  determined  by  the  analysis  of  sulfur  dioxide,  is  found  to  be  8.0175. 


318 


GENERAL   CHEMISTRY 


The  exact  atomic  weight  is  evidently  four  times  the  equivalent,  since 
this  will  give  a  number  approximating  32.  Hence  the  true  atomic 
weight  of  sulfur  is  4  x  8.0175,  or  32.07. 

The  molecular  weights  of  the  elements.  If  the  molecular  weights 
of  the  elements  themselves  are  determined  by  measurement  of  their 
vapor  densities  at  temperatures  above  their  boiling  points,  very  inter- 
esting results  are  obtained,  as  is  shown  in  the  following  table  : 

THE  MOLECULAR  WEIGHTS  OF  SOME  ELEMENTS 


ELEMENT 

TEMPERATURE 

MOLECULAR 
WEIGHT 

FORMULA 

Sodium     .      .     . 

Red  heat 

25.4 

Na  =    23.0 

Potassium     .     . 

Red  heat 

37.6 

K     =    39.1 

Zinc     .... 

1740° 

76.4 

Zn  =    65.37 

Cadmium 

1040° 

114.0 

Cd  =  112.4 

Mercury  .     .     . 

448° 

198.5 

Hg  =  200.6 

Oxygen    .     .     . 
Nitrogen  .     .     . 
Hydrogen 
Chlorine  . 

Up  to  1690° 
Up  to  1690° 
Up  to  1690° 
200° 

32.0 
28.08 
2.005 
70.9 

02  -    32. 

N2  =    28.02 
H2  =      2.016 
C12  =    70.92 

Iodine 
Iodine 

448° 
1700° 

254.8 
127. 

I2     =  253.84 
I      =  126.92 

Thallium      .     . 

1730° 

412.4 

T12  -  408. 

Ozone  .... 
Phosphorus  .     . 
Phosphorus  .     . 
Sulfur       .     .     . 

0° 
313° 
1700° 
193° 

47.9 
128.0 
91.2 
251. 

O3   =    48. 
P4    =124.16 

*4     +P2 

S8    =256.56 

Sulfur       .     .     . 

1719° 

63.6 

S2    =    64.14 

From  this  table  it  appears  that  the  metallic  elements  have  molecular 
weights  which  are  identical  with  their  atomic  weights  as  deduced 
from  their  compounds.  This  is  true  of  all  the  metals  whose  molecular 
weights  have  been  determined  by  vapor-density  methods,  with  the 
exception  of  thallium.  Judging  by  other  methods  of  measurement  it 
is  also  true  of  the  gases  constituting  Group  0  in  the  periodic  table, 
all  of  which  appear  to  be  monatomic.  Aside  from  these,  the  elements 
which  usually  occur  in  the  gaseous  state,  as  well  as  many  others, 
have  molecular  weights  double  their  atomic  weights,  the  molecule 
consisting  of  two  atoms.  As  the  temperature  is  raised,  many  of  these 
molecules,  as  with  iodine,  show  a  marked  tendency  to  break  down 
into  single  atoms,  setting  up  a  state  of  equilibrium  between  the 


MOLECULAR  WEIGHTS 


319 


molecule  and  the  atoms.  Ozone  appears  to  consist  of  three  atoms  of 
oxygen.  Phosphorus,  as  well  as  arsenic,  has  four  atoms  in  the  mole- 
cule at  ordinary  temperatures,  but  at  high  temperatures  approaches  a 
diatomic  molecule.  Sulfur  molecules  consist  of  eight  atoms  at  low 
temperatures  (measured  under  very  small  pressures),  but  at  high 
temperatures  these  decompose  into  diatomic  molecules.  At  inter- 
mediate temperatures  there  is  equilibrium  between  the  two  forms. 

The  law  of  Dulong  and  Petit.  As  early  as  1819  Dulong  and  Petit  discovered  a 
relationship  between  the  atomic  weights  of  solid  elementary  substances  and  their 
specific  heats,  which  has  been  of  much  assistance  in  fixing  upon  the  multiple 
of  the  combining  weight  which  correctly  represents  the  atomic  weight.  Their 
generalization  was  of  special  service  at  a  time  before  it  was  possible  to  deter- 
mine the  vapor  density  at  high  temperatures. 

These  investigators  found  that  the  atomic  weight  multiplied  by  the  specific 
heat  gives  approximately  a  constant  whose  value  is  about  6.25.  This  is  called 
the  atomic  heat  of  the  elements.  By  the  specific  heat  is  meant  the  quantity  of 
heat  required  to  raise  the  temperature  of  a  gram  of  the  solid  substance  one 
degree.  Evidently  the  approximate  atomic  weight  of  an  element  will  be  given 

by  the  equation : 

Atomic  weight  =  6.2o  -^-  sp.  ht. 

Many  more  recent  researches  have  been  carried  out  on  this  subject,  notably  by 
Regnault,  and  the  following  table  gives  some  of  the  values  accepted  at  the  pres- 
ent time.  In  the  case  of  some  of  the  elements,  notably  boron,  carbon,  silicon, 
and  glucinum,  the  value  obtained  deviates  widely  from  the  average  under  ordi- 
nary conditions,  but  approaches  the  normal  value  at  high  temperatures. 

TABLE  OF  SPECIFIC   HEATS 


ELEMENT 

ATOMIC 
WEIGHT 

SPECIFIC 
HEAT 

ATOMIC 
HEAT 

Lithium       

6.94 

0.941 

6.53 

Sodium  .           .     . 

23.00 

0.293 

6.74 

Magnesium      .... 
Aluminium      .... 
Phosphorus      .... 
Sulfur 

24.32 
27.1 
31.04 
32.07 

0.245 
0.214 
0.202 
0^03 

5.95 
5.80 
6.26 
6.51 

Potassium  

39.10 

0.166 

6.49 

55.84 

0.112 

6.26 

63.57 

0.095 

6.04 

Zinc              .     .           .     . 

65.37 

0.093 

6.07 

Silver      

107.88 
195.2 

0.057 
0.0325 

6.15 
6.34 

Gold 

197.2 

0.0324 

6.40 

Mercury 

200.6 

0.0333 

6.66 

Lead  

207.1 

0.0315 

6.52 

320  GENERAL  CHEMISTRY 

Some  applications  to  chemical  calculations.  The  relations  brought 
out  in  connection  with  Avogadro's  hypothesis  may  be  turned  to 
account  in  many  calculations  involving  gas  volumes. 

1.  Volume  changes  in  reactions.  In  any  reaction  which  involves  gase- 
ous products  the  volume  changes  may  be  at  once  noted.  For  example, 
when  carbon  monoxide  burns,  we  have  the  equation 

2CO  +  O2=2CO2 

Since  equal  volumes  contain  the  same  number  of  molecules,  it  is 
evident  that  the  converse  is  true:  namely,  when  equal  numbers  of 
molecules  are  present,  the  volumes  are  equal.  The  double  formula 
weight  of  carbon  monoxide  will  have  the  same  number  of  molecules 
as  the  double  formula  weight  of  the  dioxide,  and  each  of  these  will  be 
twice  the  volume  of  the  single  formula  weight  of  oxygen.  The  equa- 
tion may  therefore  be  read  :  2  x  22.4  1.  of  carbon  monoxide  +  22.4  1.  of 
oxygen  produce  2  x  22.4  1.  of  carbon  dioxide.  As  a  result  of  the  combus- 
tion there  is  therefore  a  shrinkage  of  one  third  of  the  original  volume. 
In  the  combustion  of  marsh  gas  we  have  the  following  equation  : 


There  is  the  same  number  of  molecules  on  each  side  of  the  equation, 
and  (supposing  that  the  water  remains  as  a  vapor)  the  volumes  are 
consequently  equal.  There  is  no  change  in  volume  on  combustion. 
With  benzene  there  is  a  slight  expansion,  as  is  shown  in  the  following 
equation  :  2  Q^  +  15  Q2  =  12  CO2  +  6  H2O 

In  this  case  17  volumes  produce  18.  If  the  combustion  takes  place 
under  such  conditions  that  the  benzene  is  a  liquid  at  the  outset  and 
the  water  is  liquid  at  the  conclusion  of  the  reaction,  then  15  volumes 
produce  12,  the  volumes  of  the  two  liquids  being  so  small  as  to  be 
negligible. 

Effect  of  temperature.  All  the  statements  of  the  last  paragraph 
presuppose  that  there  is  no  change  in  the  temperature  between  the 
measurements  of  volumes.  As  a  matter  of  fact,  all  these  reactions 
evolve  much  heat,  and  since  the  specific  heat  of  gases  is  small,  the 
temperature  of  the  resulting  gases  is  very  greatly  raised.  If  the  initial 
measurements  are  made  at  0°,  each  rise  in  temperature  of  273°  will 
double  the  volume  of  the  gases  formed,  and  as  the  temperature  may 
easily  go  as  high  as  1500°,  it  will  be  seen  that  there  may  be  a  momen- 
tary expansion  to  5  or  6  times  the  calculated  volume. 


MOLECULAR  WEIGHTS  321 

2.  Volume  of  a  gas  evolved  from  a  solid.  The  properties  of  the  gram- 
molecular  volume  suggest  a  direct  method  of  calculating  the  volume 
of  a  gas  produced  in  a  reaction  without  first  calculating  its  weight. 
For  example,  take  the  equation 

CaC08  +  2  HC1  =  CaCl2  +  H2O  +  CO2 

It  is  evident  that  1  gram-molecular  weight  of  calcium  carbonate  will  pro- 
duce 1  gram-molecular  volume  of  carbon  dioxide.  Since  the  molecular 
weight  of  the  former  is,  in  round  numbers,  100,  it  follows  that  100  g. 
of  the  carbonate  will  produce  22.4  1.  of  the  dioxide,  and  the  same 
ratio  will  hold  between  any  other  quantities  in  the  reaction.  Thus,  if 
it  is  desired  to  know  the  volume  produced  from  20  g.,  it  is  sufficient 
to  note  that  this  is  ^-fa  of  the  molecular  weight,  and  that  as  a  result 
•f-fa  of  the  gram-molecular  volume  of  gas  will  be  produced,  that  is, 
_2_o_  x  22.4  =  4.48  1.  It  will  be  seen  that  it  is  not  necessary  to  know 
the  weight  of  a  liter  of  the  gas  in  order  to  make  such  calculations. 

3.  Weight  of  a  liter  of  a  gas.  It  is  often  desirable  to  know  the  weight 
of  a  liter  of  various  gases,  and  occasionally  it  is  not  convenient  to 
search  for  the  values  in  tables.    It  is  always  possible  to  obtain  an 
approximate  value  by  recalling  the  fact  that  the  molecular  weight  in 
grams  occupies  22.4  1.    If  the  formula  of  the  gas  is  known,  it  is  there- 
fore sufficient  to  divide  the  molecular  weight  by  22.4  in  order  to  de- 
termine the  weight  of  1  1.    The  weight  so  obtained  is  approximate 
only  for  the  reason  that  the  gas  laws  are  only  an  approximately  cor- 
rect statement  of  the  conduct  of  gases,  and  consequently  Avogadro's 
hypothesis,  which  is  based  upon  these  laws,  cannot  be  rigidly  true. 


CHAPTER  XXIII 


FLAMES;  FUEL  GASES;  EXPLOSIONS 

Visible  combustion.  When  combustion  proceeds  rapidly,  the  heat 
liberated  is  readily  perceived,  and,  as  previously  stated,  the  reaction  is 
accompanied  by  light.  The  products  of  combustion  may  be  solids,  as 
in  the  case  of  metals  such  as  iron ;  or  they  may  be  liquids,  as  in  the 
case  of  hydrogen ;  or  gases,  as  with  carbon  and  sulfur.  If  the  burn- 
ing substance  is  a  solid  at  the  temperature  of  combustion,  it  may 
become  incandescent,  but  there  is  no  flame.  This  is  the  case  with  the 
combustion  of  pure  carbon  and  many  of  the  metals.  When  both  of 
the  substances  concerned  in  the  combustion  are  gases,  the  bounding 
surface  between  the  two  presents  the  appearance  known  as  a  flame, 

and  this  flame  is  usually,  though  not 
always,  distinctly  luminous.  The  fact 
that  a  flame  often  accompanies  the  com- 
bustion of  solids,  such  as  coal,  wood,  or 
a  candle,  or  of  liquids,  such  as  oils, 
does  not  contradict  the  statements  just 
made.  It  can  be  shown  that  in  each 
case  the  heat  of  combustion  produces 
vapors  from  the  burning  solid  or  liquid, 
and  these  in  turn  burn  with  a  flame. 
If  one  end  of  a  slender  glass  tube  is 
held  in  the  base  of  a  candle  flame,  as 

indicated  in  Fig.  Ill,  a  flame  is  formed  when  a  light  is  applied  at 
the  other  end,  showing  that  gases  which  can  be  drawn  off  and  ignited 
separately  are  produced  from  the  candle.  The  blue  flame  over  a  bed 
of  coals  is  the  flame  of  burning  carbon  monoxide  formed  by  the  com- 
bustion of  the  carbon. 

Fuel  gases.  Before  considering  the  structure  of  flames  it  will  be  of 
advantage  to  have  before  us  the  general  chemical  characteristics  of  the 
gases  whose  combustion  gives  rise  to  the  most  familiar  examples  of 
flames.  A  number  of  varieties  of  gases  are  now  employed  as  sources 
of  heat,  light,  and  power. 


FIG.  Ill 


FLAMES;   FUEL  GASES;  EXPLOSIONS 


323 


1.  Coal  gas.  It  has  been  known  for  several  centuries  that  when  soft, 
or  bituminous,  coal  is  heated  out  of  contact  with  air,  combustible 
gases  are  formed,  and  gas  obtained  in  this  way  was  used  in  street 
lighting  in  London  and  Paris  a  hundred  years  ago. 

The  manufacture  of  coal  gas  is  represented  in  a  diagrammatic  way  in  Fig.  112. 
A  represents  one  of  the  closed  retorts  in  which  the  coal  is  placed,  and  which  is 
heated  by  the  fire  below.  A  number  of  these  are  placed  in  parallel  rows,  each 
being  furnished  with  a  delivery  pipe,  from  which  the  gas  bubbles  into  the  tarry 
liquids  which  collect  in  the  hydraulic  main  B,  running  above  the  retorts.  In 
this  large  pipe  are  deposited  most  of  the  solid  and  liquid  products  formed  in 
distillation,  constituting  the  viscous  mass  known  as  coal  tar.  The  partially 
purified  gas  then  passes  into  a  series  of  pipes-  C,  in  which  it  is  cooled  and  further 
separated  from  tar.  In  the  scrubber  D  it  passes  through  a  column  of  loose  coke 


FIG.  112 


over  which  water  is  sprayed,  where  it  is  still  further  cooled  and  to  some  extent 
purified  from  soluble  gases,  such  as  hydrogen  sulfide  and  ammonia.  In  the 
purifier  E  it  passes  over  a  bed  of  lime  or  iron  oxide,  which  removes  the  remain- 
der of  the  sulfur  compounds,  and  from  this  it  enters  the  large  gas  holder  F, 
from  which  it  is  distributed  to  consumers. 

The  great  bulk  of  the  carbon  remains  in  the  retort  as  coke  and  retort  carbon. 
The  yield  of  gas,  tar,  and  soluble  materials  depends  upon  many  factors,  such 
as  the  composition  of  the  coal,  the  temperature  employed,  and  the  rate  of  heat- 
ing. One  ton  of  good  gas  coal  yields  approximately  10,000  cu.  ft.  of  gas,  1400  Ib. 
of  coke,  120  Ib.  of  tar,  and  20  gal.  of  ammoniacal  liquor. 

2.  Water  gas.  Water  gas  is  essentially  a  mixture  of  carbon  monoxide 
and  hydrogen.  It  is  manufactured  by  passing  superheated  steam  over 
very  hot  anthracite  coal  or  coke,  the  chief  reactions  being  expressed 
in  the  following  equations: 

C  -f  H20  =  CO  +  H2  -  26,990  cal. 
CO2  +  C  =  2  CO  -  37,230  cal. 

The  industrial  process  is  intermittent.  The  fuel  is  burned  with  a  forced  draft 
in  a  suitable  furnace  until  it  is  very  hot.  The  air  is  then  shut  off  and  the  steam 
turned  on  until  the  temperature  falls  to  about  1000°.  The  process  is  then 


324  GENERAL   CHEMISTRY 

reversed.  The  fall  in  temperature  is  rapid,  partly  owing  to  radiation  and  to  the 
cooling  occasioned  by  the  steam,  but  largely  because  of  the  endothermic  char- 
acter of  the  reactions  which  take  place.  The  gas  so  formed  contains  all  the 
nitrogen  which  was  in  the  furnace  when  the  steam  was  admitted. 

Water  gas  burns  with  a  pale-blue,  nonluminous  flame.  It  is  very 
poisonous  and  has  no  odor.  To  make  it  suitable  for  illumination  in 
an  ordinary  burner,  as  well  as  to  give  it  an  odor  and  so  make  it  safer, 
it  must  be  enriched  with  hydrocarbons  called  illuminants.  This  is 
accomplished  by  passing  the  gas  through  a  furnace  filled  with  hot  fire 
brick  upon  which  crude  petroleum  is  sprayed.  The  petroleum  oils  are 
decomposed  (cracked)  into  simpler  gaseous  bodies,  the  most  important 
of  which  are  methane,  acetylene,  and  ethylene.  Coal  gas  is  sometimes 
enriched  in  a  similar  way  by  adding  petroleum  to  the  coal  in  the  retorts. 

Gas  mantles.  Instead  of  depending  upon  illuminants  added  to  the 
gas,  it  is  much  better  to  suspend  a  gauze  of  suitable  material  around 
the  nonluminous  hot  flame,  the  incandescence  of  the  gauze  furnishing 
the  light.  The  various  physical  and  chemical  requirements  on  the 
part  of  such  materials  are  very  difficult  to  meet,  but  it  has  been  found 
that  a  mixture  of  the  oxides  of  thorium  and  cerium  in  the  ratio  of 
99  :  1  serves  admirably.  Any  deviation  from  this  ratio  decreases  the 
luminosity,  and  few  other  materials  have  any  permanent  efficiency. 

3.  Producer  gas.    This  gas  is  used  in  connection  with  many  metallur- 
gical furnace  operations,  and  as  a  fuel  for  gas  engines.    It  is  made  by 
burning  coal  under  such  conditions  that  the  product  of  combustion 
is  largely  carbon  monoxide.    The  gas  usually  consists  of  about  60  per 
cent  of  nitrogen  and  40  per  cent  of  carbon  monoxide.    It  can  be  made 
from  coal  of  a  poor  quality,  even  from  lignite,  and  as  gas  engines 
run  well  with  this  gas,  it  furnishes  the  most  economical  method  for 
utilizing  low-grade  coal  for  power. 

4.  Oil  gas.  In  countries,  such  as  the  United  States,  in  which  petro- 
leum is  found  abundantly,  some  illuminating  gas  is  made  from  it  by 
spraying  the  petroleum  upon|very  hot  fire  bricks,  as  in  the  enriching 
of  water  gas.    Pintsch  gas  is  made  in  this  way  by  cracking  petroleum 
oils,  and  is  stored  under  pressure  in  cylinders  for  lighting  railway  cars. 

5.  Natural  gas.    In  many  regions  of  the  United  States,  as  well  as 
in  other  countries,  natural  gas  is  obtained  from  wells  drilled  into  a 
stratum  holding  the  gas.    While  it  is  variable  in  composition,  it  con- 
sists largely  of  methane,  many  samples  running  as  high  as  95  per  cent 
of  this  compound.    It  burns  with  a  rather  smoky  flame  of  moderate 


FLAMES;  FUEL   GASES;  EXPLOSIONS 


325 


luminosity,  but  works  well  with  a  gas  mantle.    It  has  a  high  heat  of 
combustion,  as  shown  in  the  following  equation: 

CH4  +  2  O2  =  CO2  +  2  H2O  +  211,930  cal. 

Comparative  composition  of  gases.  The  following  figures  are  the  results  of 
analyses  of  average  samples,  but  since  each  kind  of  gas  varies  considerably  in 
composition,  the  values  are  to  be  taken  as  approximate  only.  The  nitrogen  and 
traces  of  oxygen  are  derived  from  the  air. 

COMPOSITION  OF  GASES 


OHIO 
NATURAL 
GAS 

COAL 
GAS 

WATER 
GAS 

ENRICHED 
WATER 
GAS 

PRODUCER 
GAS 

H 

00 

41.3 

52.88 

37.96 

10  90 

"2      * 

CHA 

89.5 

43.6 

2.16 

7.09 

C6H6     

9.3 

2.01 

C2H2  +  C2H4      .     .     . 
CO             .                   . 

0.3 
0.4 

3.9 
6.4 

36.80 

9.40 
32.25 

0.60 
20.10 

CO 

03 

20 

3.47 

4.73 

8  50 

N, 

0.2 

1.2 

4.69 

3.96 

59.90 

O, 

0.0 

0.3 

0.60 

Other  hydrocarbons 

0.0 

1.5 

1.80 

Relation  of  the  two  gases  to  the  flame.  The  gas  issuing  from  the 
burner  is  said  to  undergo  combustion,  while  that  one  which  constitutes 
the  atmosphere  about  the  flame  is  said  to 
support  combustion.  These  terms  are  en- 
tirely conventional,  since  the  relation  of  the 
two  gases  may  be  reversed  without  greatly 
altering  the  appearance  of  the  flame. 

c=c 
Fig.  113  illustrates  a  convenient   apparatus  for 

demonstrating  this  fact.   A  wide  lamp  chimney  A  is 

covered  with  a  piece  of  asbestos  board  B,  which  has 

a  hole  in  the  center  about  as  large  as  a  dime.    A 

straight  tube  C,  about  1  cm.  wide,  and  also  a  smaller 

tube  D  connected  with  the  gas  supply  pass  through 

a  cork  at  the  bottom.  If  the  hole  in  B  is  closed  (by 

a   piece  of  asbestos  board)  while  gas  is  admitted 

through  D,  the  excess  gas  escapes  downward  through 

C,  where  it  may  be  lighted.    The  hole  in  B  is 'now 

opened,  the  flame  ascending  to  the  top  of  the  tube  C.  This  flame  is  produced  by 

air,  drawn  up  through  C,  burning  in  an  atmosphere  of  coal  gas.    Finally,  the 

excess  of  coal  gas  may  be  ignited  at  B,  where  it  will  burn  in  air,  the  two  flames 

being  very  similar. 


FIG.  113 


326 


GENERAL  CHEMISTRY 


FIG.  114 


Structure  of  a  flame.  The  structure  of  a  flame  can  be  studied  to 
the  best  advantage  when  the  combustible  gas  issues  from  a  round 
tube  into  an  atmosphere  of  the  gas  supporting  combustion  (usually 
the  air),  as  is  the  case  with  an  ordinary  Bunsen  burner 
(Fig.  114).  Under  these  conditions  the  flame  is  conical 
in  outline. 

Simple  flames.  When  the  chemical  action  taking  place 
in  the  combustion  is  the  mere  union  of  two  gases,  as  is 
true  in  the  union  of  hydrogen  or  carbon  monoxide  with 
oxygen,  or  hydrogen  with  chlorine,  the  structure  of  the 
flame  is  very  simple.  It  consists  of  two  superimposed 
cones  of  different  altitudes.  The  inner  one  may  be  shown 
to  be  merely  unchanged  cold  gas,  and  is  therefore  not  a 
real  part  of  the  flame.  A  match  head  suspended  in  this 
region  (Fig.  114)  before  lighting  the  gas  is  not  ignited 
by  the  flame  around  it. 

Complex  flames.  In  the  burning  of  hydrocarbons,  as 
well  as  of  many  other  gases,  the  flame  is  more  complex, 
and  as  many  as  four  distinct  cones  may  be  seen  (Fig.  115). 
The  innermost  one  (A)  is  really  not  a  part  of  the  flame,  being  formed 
of  gas  not  yet  brought  to  the  point  of  combustion.  If  a 
Bunsen  burner  is  employed,  with  the  ring  at  the  base 
turned  to  admit  plenty  of  air,  the  second  cone  (^)  is 
sharply  defined  and  is  bluish  green  in  color.  If  the 
burner  tube  is  wide,  or  too  much  air  is  admitted,  the 
rate  of  combustion  in  this  cone  may  exceed  the  rate  of 
flow  of  the  gas,  in  which  case  the  cone  will  travel  down 
the  tube  and  burn  at  the  base,  or  strike  back.  As  the 
air  is  shut  off  it  will  be  seen  that  a  luminous  spot  ap- 
pears at  the  apex  of  this  cone,  which  gradually  takes  the 
form  of  a  cone  ((7)  quite  covering  the  inner  one  and 
brightly  luminous  over  all  its  surface.  Finally,  if  some 
object  is  held  so  as  to  intercept  the  light  from  this  region, 
it  will  be  seen  that  there  is  a  fourth  cone  (Z>),  which  is 
only  faintly  luminous. 

Cause  of  the  cones.    Since  the  gases  which  give  rise  to 
multiplication  of  cones  on  combustion  are  always  chemical 
compounds,  such  as  hydrocarbons,  it  would  appear  probable  that  the 
cones  are  due  to  successive  chemical  reactions  taking  place  in  different 


—B 


— A 


FIG.  115 


FLAMES;  FUEL   GASES;  EXPLOSIONS  327 

regions.  Smithells  devised  a  very  simple  method  for  separating  the 
two  principal  cones  and  ascertaining  the  character  of  the  combustion 
in  each  one.  The  essentials  of  his  apparatus  are  shown  in  Fig.  116. 

The  tube  A  of  an  ordinary  Bunsen  burner  is  extended  in  length  by  a  glass 
tube  B  of  slightly  greater  diameter.  A  wider  tube  C,  of  about  2  cm.  diameter 
and  provided  with  a  side  tube  D,  surrounds  this  tube,  to  which  it  is  connected  by 
a  cork  E,  which  slips  readily  on  the  smaller  tube.  The  wide  tube  Cis  slipped  down 
until  the  ends  of  the  two  tubes  are  even;  the  gas  is 
turned  on  and  lighted  at  the  top  of  these  tubes,  where 
it  burns  with  a  double  cone.  Air  is  admitted  at  the 
base  of  the  burner  until  the  inner  cone  is  sharply  de- 
nned and  is  bluish  green  in  color.  The  outer  tube  is 
then  pushed  up,  carrying  the  outer  cone  with  it  and 
leaving  the  inner  one  upon  the  rim  of  the  smaller  tube,  ^ 

as  shown  in  the  figure.  The  two  cones  are  widely  sepa- 
rated in  this  way,  and  the  space  between  them  becomes 
filled  with  the  gases  formed  by  the  combustion  in  the 
inner  cone.  These  can  be  drawn  off  through  the  side 
tube  D  and  analyzed. 

In  this  way  Smithells  showed  that  in  the 
inner  cone  the  original  hydrocarbons  burn  to 
form  carbon  monoxide  and  hydrogen,  together 

with  some  unsaturated  hydrocarbons,  chiefly  acetylene  and  ethylene. 
In  the  second,  or  luminous,  cone  the  hydrogen  and  carbon  monoxide 
are  in  part  burned  by  the  oxygen  supply  in  the  atmosphere,  while  the 
illuminants,  acetylene  and  ethylene,  are  decomposed  into  carbon  and 
hydrogen,  the  separated  carbon  becoming  brilliantly  incandescent.  It 
quickly  undergoes  combustion,  however,  and  the  outer  edge  of  the 
luminous  flame  marks  its  disappearance.  The  outside,  faint  mantle  is 
the  region  in  which  the  combustion  of  carbon  monoxide  and  hydrogen 
is  completed. 

Causes  affecting  luminosity.  While  the  process  just  sketched  ac- 
counts for  the  luminosity  of  flames  in  a  general  way,  there  are  evidently 
a  number  of  other  factors  which  must  be  taken  into  account.  Some 
flames,  such  as  that  of  burning  ammonia,  are  luminous,  though  there 
is  no  solid  incandescent  product  formed  during  combustion.  It  is 
possible  that,  in  the  decomposition  into  elements  at  a  high  temperature, 
endothermic  bodies  may  give  up  some  of  their  energy  directly  as  light 
rather  than  as  heat. 

The  temperature  of  the  gases  before  combustion  also  affects  lumi- 
nosity. The  nonluminous  flame  of  a  Bunsen  burner  becomes  somewhat 


328 


GENERAL  CHEMISTRY 


FIG.  117 


luminous  when  the  tube  of  the  burner  is  strongly  heated.  When  the 
gas  or  the  flame  itself  is  cooled,  the  luminosity  diminishes,  as  may  be 
seen  by  bringing  a  large  mass  of  cold  metal,  such  as  a  flatiron,  close 
to  a  luminous  flame.  A  loose  spiral  of  heavy  cop- 
per wire  brought  down  over  a  luminous  Bunsen 
flame  acts  in  the  same  way  (Fig.  117). 

The  concentration  of  the  gases  is  also  an  impor- 
tant factor,  the  luminosity  being  greater  as  the 
concentration  increases.  Consequently,  pressure 
increases  luminosity,  as  is  shown  by  the  fact  that 
hydrogen  under  pressure  burns  in  oxygen  (also 
under  pressure)  with  a  luminous  flame.  On  the 
other  hand,  dilution  with  an  indifferent  gas,  espe- 
cially if  it  is  cold,  greatly  reduces  luminosity. 
Carbon  dioxide,  nitrogen,  or  even  air,  admitted  at  the  base  of  a 
luminous  flame,  destroys  the  luminosity.  Such  gases  evidently  act  by 
diminishing  the  rapidity  of  combustion  and  conse- 
quently the  heat  per  unit  of  time,  and  by  absorbing 
heat  and  so  reducing  the  temperature  still  farther. 
The  temperature  of  flames.  The  actual  temper- 
ature which  can  be  realized  in  an  ordinary  flame 
obviously  depends  upon  many  conditions,  such  as 
the  composition  of  the  gas,  its  pressure,  tempera- 
ture, and  rate  of  flow,  and  the  method  of  supply- 
ing the  air.  Even  in  an  ordinary  Bunsen  flame 
burning  under  favorable  conditions  it  is  very 
difficult  to  determine  the  maximum  temperature 
attained.  The  actual  region  of  great  heat  is  very 
limited,  as  the  burning  zones  are  very  thin.  The 
temperature  in  different  parts  of  the  flame  is  very 
different,  and  any  object  placed  in  the  flame,  for 
determining  its  temperature,  cuts  across  many 
different  regions  and  is  unequally  heated.  Evi- 
dently the  temperature  is  much  higher  than  that 
recorded  by  a  body  in  the  flame,  since  the  specific 
heat  of  solids  is  so  much  greater  than  that  of  gases. 
Under  exceptional  conditions  it  has  been  found  possible  to  melt  a  very 
fine  platinum  wire  in  a  good  Bunsen  flame  so  that  a  temperature  of 
1755°  is  surely  reached.  The  accompanying  diagram  (Fig.  118}  gives 


-\-~1540' 


1550° 


1570- 


1450- 


-350" 
-300° 


FIG.  118 


FLAMES;   FUEL  GASES;   EXPLOSIONS  329 

a  rough  estimate  of  the  probable  temperature  in  various  parts  of  a 
good  nonluminous  Bunsen  flame. 

Reducing  and  oxidizing  flames.  Since  the  region  just  below  the 
luminous  cone  is  very  hot  and  contains  the  reducing  gases  hydrogen 
and  carbon  monoxide,  a  substance  such  as  a  metallic  oxide,  placed  in 
this  region,  will  undergo  reduction,  provided  it  can  be  reduced  by 
such  hot  gases.  This  region  is  therefore  called  the  reducing  region, 
and  a  body  heated  in  this  way  is  said  to  be  heated  in  the  reducing 
flame.  At  the  apex  of  the  flame  there  are  no  reducing  gases,  but  it 
is  very  hot  and  air  is  abundant;  consequently  a  substance  which  is 
rather  readily  oxidized  will  undergo  oxidation  if  heated  in  this  region. 
This  is  called  the  oxidizing  flame. 

Explosions.  An  explosion  is  caused  by  the  sudden  change  in  the 
volume  of  gases  following  chemical  reaction  or  as  the  result  of  the 
formation  of  gases  from  liquid  or  solid  materials.  The  greater  the  vol- 
ume change,  and  the  more  rapidly  it  is  produced,  the  more  violent 
the  explosion. 

The  equation  of  the  reaction  does  not  always  supply  the  informa- 
tion necessary  for  predicting  an  explosion.  Thus,  when  2  volumes 
of  hydrogen  and  1  of  oxygen  are  mixed  and  ignited,  2  volumes  of 
water  vapor  are  formed,  which  is  not  a  great  volume  change.  Taking 
into  account  the  heat  evolved,  however,  the  complete  equation  is  as 
follows : 

2  H2  +  O2  =  2  H2O  +  116,138  cal. 

If  we  remember  Gay-Lussac's  law  of  gas  expansion,  and  recall  the 
fact  that  the  specific  heats  of  gases  are  very  small,  we  shall  see  that 
the  heat  of  reaction  may  lead  to  an  expansion  of  5  or  6  volumes, 
especially  when  the  reaction  is  very  rapid. 

Explosive  mixtures.  A  second  fact  not  indicated  by  the  equation 
of  a  reaction  is  that  explosion  of  a  gas  mixture  will  not  occur  unless 
the  mixture  falls  between  certain  limits  in  percentage  composition. 
Thus,  if  hydrogen  and  air  are  mixed,  the  resulting  mixture  is  not 
explosive  unless  the  volume  percentage  of  hydrogen  is  above  5  per 
cent  and  below  72  per  cent.  In  mixtures  outside  of  these  limits  the 
combination  is  so  slow,  the  heat  is  so  largely  absorbed  by  the  excess 
gases  present,  and  the  volume  change  takes  place  so  slowly  and  is  so 
small  a  fraction  of  the  total  volume  that  no  explosion  occurs.  The 
following  table  gives  both  the  lower  and  the  upper  limits  of  explosive 


330 


GENERAL  CHEMISTRY 


mixtures  of  several  gases  with  air,  expressed  in  volume  percentages. 
The  values  are  to  be  regarded  as  roughly  approximate. 


EXPLOSIVE  MIXTURES 


GAS 

VOLUME  PERCENTAGE 
AT  LOWER  LIMIT 

VOIAJME  PERCENTAGE 
AT  HIGHER  LIMIT 

Hydrogen 

5 

72 

Methane    

5 

13 

Carbon  monoxide    .     ,  ..    . 
Acetylene  

13 
3 

75 

82 

Water  gas       .     .     .     .'    .  ' 

Coal  gas                    .     '.'•  '  'v'1 

9 
6 

55 
29 

Mine  explosions.    In  many  coal  mines  methane  collects  at  times  and 
is  called  fire  damp.    It  forms  an  explosive  mixture  with  air  when  the 
percentage  of  methane  rises  above  5  per  cent, 
the  equation  for  the  reaction  being  as  follows : 

CH4  +  2  02  =  C02  +  2  H20  +  211,930  cal. 

The  expansion  in  this  case  is  entirely  due  to 
the  heat.  In  the  reaction  the  oxygen  of  the 
air  is  decreased  and  the  carbon  dioxide  is  in- 
creased to  such  a  point  that  the  air  will  no 
longer  support  respiration.  The  gases  result- 
ing from  the  explosion  are  called  choke  damp, 
and  often  suffocate  the  miners. 

Safety  lamp.  Fortunately  the  ignition  point 
of  fire  damp  is  high  and  its  flame  may  be  extin- 
guished by  cooling.  In  1815  Sir  Humphry 

Davy  invented  a  miner's  lamp,  based  on  this 

principle,  in  which  the  usual  chimney  of  a  lan- 
tern is  replaced  by  a  wire  gauze  (Fig.  119).  An 

explosion  flame  starting  at  the  wick  is  so  cooled 

by  the  metal  wire  that  ignition  ceases  and  the 

explosion  is  confined  to  the  interior  of  the  lamp. 

The  principle  may  be  demonstrated  by  holding 

a  wire  gauze  a  few  inches   above   a  Bunsen 

flame  parallel  with  the  table  (Fig.  120).  When 

the  gas  is  turned  on  and  a  light  applied  above  the  gauze,  the  resulting 

flame  rests  upon  the  gauze,  but  does  not  pass  through  it  to  the  burner. 


FIG.  119 


FIG.  120 


FLAMES;   FUEL  GASES;  EXPLOSIONS  331 

Dust  explosions.  In  dry  mines  great  quantities  of  fine  coal  dust 
collect  if  the  mine  is  not  kept  damp  by  spraying.  A  blast  may  blow 
this  into  the  air,  and  the  result  may  be  an  explosion  due  to  the  union 
of  carbon  and  oxygen  : 

2  C  +  O2  =  2  CO  +  58,000  cal. 

Moreover,  the  carbon  monoxide,  called  afterdamp,  is  also  explosive 
when  mixed  with  air,  and  is  very  poisonous.  Rescuers  often  carry 
live  birds  with  them,  as  these  are  extremely  sensitive  to  the  poisonous 
effects  of  carbon  monoxide,  and  their  death  warns  the  rescuers  of 
their  own  peril. 

In  a  similar  way  severe  explosions  have  resulted  in  flour  mills  and 
woodworking  plants  from  the  fine  combustible  dust  floating  in  the  air, 
so  that  in  modern  mills  all  machinery  producing  such  dust  is  covered 
by  a  hoqd  in  which  a  strong  draft  is  maintained. 

Explosives.  Manufactured  explosives  are  of  two  general  classes. 

1.  Low  explosives.  The  first  is  represented  by  ordinary  gunpowder, 
and  is  a  more  or  less  intimate  mixture  of  solids  which  combine  rapidly 
when  ignited  and  produce  gaseous  products.  The  composition  and 
reaction  of  gunpowder  is  expressed  in  a  general  way  by  the  following 
equation  :  +  g  =  R 


2.  High  explosives.  In  the  other  class,  known  as  high  explosives,  the 
effect  is  produced  by  the  spontaneous  decomposition  of  endothermic 
substances.  Most  of  these  are  nitrates  of  organic  substances  and  in- 
clude nitroglycerin  (C8H5(NO8)3)  and  nitrocellulose  (C12HMO4(NO8)6). 
High  explosives,  when  ignited,  burn  rapidly  but,  as  a  rule,  not  with 
explosive  violence.  The  explosion  is  brought  about  by  a  shock  pro- 
duced by  an  explosive  cartridge  ignited  by  a  fuse.  Owing  to  the  almost 
instantaneous  character  of  the  explosion,  an  enormous  gas  pressure  is 
suddenly  produced  in  all  directions,  which  will  shatter  a  solid  rock  on 
which  the  explosive  is  placed.  This  gives  rise  to  the  popular  idea  that 
dynamite  explodes  downward,  which  of  course  is  not  true. 


CHAPTER  XXIY 

THERMOCHEMISTRY 

In  Chapter  I  it  was  pointed  out  that  there  are  certain  respects  in 
which  a  parallel  exists  between  the  potential  energy  of  bodies  and 
the  chemical  energy  which  one  substance  has  for  another.  The  same 
parallel  may  be  employed  to  illustrate  another  aspect  of  chemical 
energy  which  must  now  be  considered. 

Energy  liberated  in  stages.  When  a  stone  falls  from  the  top  of  a 
building  to  the  pavement  below,  its  potential  energy  is  converted  into 
kinetic  energy  and,  on  striking  the  pavement,  into  mechanical  effects 
and  heat.  If  the  stone  should  chance  to  come  to  rest  near  the  brink 
of  a  deep  well,  a  very  little  effort  would  serve  to  push  it  over  the 
brink,  when  it  would  fall  farther,  with  the  transformation  of  an  addi- 
tional quantity  of  potential  energy  into  other  forms.  Evidently  the 
process  may  continue  until  the  body  reaches'  the  center  of  the  earth, 
and  its  total  potential  energy  at  any  point  may  be  calculated  as  equal 
to  the  work  done  upon  it  in  raising  it  from  the  center  of  the  earth 
to  the  position  in  question. 

In  a  somewhat  parallel  way  two  elements  may  be  regarded  as  each 
possessing  a  certain  quantity  of  chemical  energy.  When  they  unite, 
a  certain  portion  of  this  is  given  off  as  heat,  electrical  energy,  or  light, 
but  the  compound  so  formed  still  possesses  the  ability  to  unite  with 
other  elements  and  compounds,  liberating  an  additional  quantity  of 
energy.  This  is  illustrated  in  the  following  reactions: 

Zn  +  S  (rhombic)  =  ZnS  +  39,570  cal. 

ZnS  +  2O2  =  ZnSO4  +  190,500  cal. 
ZnSO4  +  7  H2O  -  ZnSO4  •  7  H2O  +  22,690  cal. 

Total,  free,  and  bound  energy.  It  is  quite  possible  to  determine  the 
energy  liberated  in  each  one  of  these  stages,  just  as  we  may  determine 
the  loss  of  potential  energy  at  each  stage  of  a  stone's  fall.  We  have 
no  way,  however,  of  getting  at  the  total  chemical  energy  in  an  element, 
for  we  can  never  examine  it  when  it  has  lost  all  its  power  of  combina- 
tion, nor  have  we  any  theoretical  basis  for  calculating  it,  as  we  have 

332 


THERMOCHEMISTRY  333 

with  the  potential  energy  of  a  stone.  All  we  can  do  is  to  deter- 
mine the  difference  in  energy  between  two  different  stages  of  combi- 
nation. It  is  clear  that  the  heat  of  a  reaction  gives  no  information  as 
to  the  total  energy  in  the  reacting  substances,  but  only  the  change 
in  chemical  energy  which  they  experience  on  combining.  This  latter 
quantity  is  usually  spoken  of  as  the  free  energy.  The  remaining 
energy,  which  we  cannot  measure,  is  called  the  bound  energy.  There 
is  good  reason  for  believing  that  the  latter  is  very  much  greater  than 
the  former. 

Chemical  energy  a  relative  quantity.  There  is  also  this  important 
distinction  between  gravitational  and  chemical  energy.  All  materials, 
independently  of  then:  chemical  composition,  are  attracted  equally, 
mass  for  mass,  by  the  earth.  Every  element,  on  the  other  hand,  has 
its  own  peculiar  affinity  for  every  other  element,  and  there  is  no 
way  of  determining  this  except  by  measurement.  We  cannot  state 
what  energy  the  element  carbon  possesses,  but  can  merely  tell  how 
much  it  evolves  when  it  combines  with  hydrogen  or  with  oxygen  or 
with  sulfur. 

Heat  of  reaction  not  proportional  to  chemical  affinity.  While  the  heat 
of  a  reaction  often  gives  us  a  clear  idea  as  to  the  intensity  of  the 
chemical  affinity  between  two  substances,  this  is  not  always  so,  and 
certain  facts  must  be  remembered  in  drawing  inferences  from  it. 

1.  Heat  effect  accompanying  changes  in  state.  It  frequently  happens 
that  changes  in  state  or  crystalline  form  accompany  chemical  reactions, 
and  these  modify  to  quite  an  extent  the  total  heat  change  in  the 
reaction.  For  example,  when  solid  rhombic  sulfur  is  burned,  we  have 
the  equation  g  (rhombic)  +  ^  =  s^  +  71)0go  cal 

When  liquid  sulfur  is  burned,  the  equation  is 

S  (liquid)  +  O2  =  SO2  +  71,380  cal. 

The  difference,  300  cal.,  represents  the  heat  of  fusion  of  the  sulfur. 

In  like  manner,  when  oxygen  and  hydrogen  combine  to  form  steam 
at  100°,  the  equation  is 

2  H2  +  O2  =  2  H2O  +  116,138  cal. 

If  the  heat  is  measured  under  such  conditions  that  the  resulting  water 
is  liquid  at  20°,  the  equation  becomes 

2  H2  +  O2  =  2  H2O  -1-  136,684  cal. 


334  GENERAL  CHEMISTRY 

2.  Heat  of  endothermic  compounds.   Many  substances  are  endothermic 
in  character,  evolving  heat  on  decomposition.   If  the  heat  of  formation 
were  to  be  taken  as  a  direct  measure  of  chemical  affinity,  there  should 
be  less  than  no  affinity  between  the  elements  of  such  a  compound, 
and  therefore  no  force  to  hold  them  together ;  yet  many  endothermic 
bodies,  such  as  carbon  disulfide  and  acetylene,  are  fairly  stable  under 
ordinary  conditions. 

3.  Chemical  reactions  dependent  upon  secondary  changes.  The  heat  due  to  causes  aside 
from  the  reaction  itself  may  greatly  modify  the  course  of  the  reaction  and  even 
reverse  its  direction.    This  may  be  seen  from  a  study  of  an  example.    Hydrogen 
and  iodine  combine  to  form  gaseous  hydrogen  iodide,  with  heat  absorption : 

H2  +  I2  =  2  HI  -  12,072  cal. 

Hydrogen  and  sulfur  combine  to  form  hydrogen  sulfide,  with  heat  evolution : 
H2  +  S  =  H2S  +  2,730  cal. 

We  should  therefore  expect  sulfur  to  decompose  hydrogen  iodide,  with  heat 
evolution,  which  it  does : 

2  HI  +  S  =  H2S  +  I2  +  14,802  cal. 

In  solution,  on  the  other  hand,  this  action  is  reversed,  for  the  heat  of  solution  of 
all  the  factors  concerned  comes  into  the  account : 

H2  +  I2  (+  water)  =  2  HI  +  26,342  cal. 
Ha  +  S  (  +  water)  =  H2S  +  7,290  cal. 

The  action  of  iodine  upon  hydrogen  sulfide  in  solution  may  therefore  be  ex- 
pressed by  the  equation 

H2S  +  Ia  =  2  HI  +  S  +  19,052  cal. 
This  value,  19,052  cal.,  is  the  algebraic  sum  of  the  following : 

Heat  absorbed  in  formation  of  2  HI  (gaseous)        ...  —  12,072  cal. 

Heat  absorbed  in  decomposing  H2S  (in  solution)        .     .  —    7,290  cal. 

Heat  evolved  in  solution  of  2  HI  (gaseous) +  38,414  cal. 

Total  heat  change  in  the  reaction +  19,052  cal. 

The  measurement  of  heat  of  reaction.  Two  general  types  of  calo- 
rimeters are  employed  in  measurements  of  the  heat  of  reactions.  In 
reactions  taking  place  in  solution  the  open  calorimeter  described  on 
page  7  may  be  used.  In  such  reactions  as  combustion,'  where  a  gase- 
ous substance  must  be  supplied  to  maintain  the  action,  a  bomb 
calorimeter  is  used. 

Bomb  calorimeter.  This  is  a  strong  steel  flask  lined  with  platinum  or  porcelain 
and  provided  with  a  tight-fitting  screw  cap  (Fig.  121).  In  determining  the  heat 
of  combustion  a  weighed  sample  of  the  substance  is  placed  on  the  capsule  A, 


THERMOCHEMISTRY 


335 


D 


oxygen  is  admitted  through  the  tube  B  until  the  pressure  in  the  bomb  is  about 
20  atmospheres,  and  the  bomb  is  then  closed  and  placed  in  an  open  calorimeter. 
The  charge  is  ignited  by  passing  an 
electric  current  through  the  fine  iron 
fuse-wire  C  stretched  above  the  charge. 
The  wire  is  melted  and  the  red-hot  drop 
of  burning  metal  falls  upon  the  charge, 
igniting  it.  The  heat  given  off  during 
combustion  is  measured  by  the  rise  in 
temperature  of  the  water  surrounding 
the  bomb,  which  is  stirred  by  the  stirrer 
D.  A  preliminary  experiment  must  be 
made  upon  a  weighed  charge  of  a  sub- 
stance whose  heat  of  combination  is 
known  (such  as  cane  sugar),  to  deter- 
mine the  heat  absorbed  by  the  bomb, 
together  with  that  due  to  the  melting 
and  combustion  of  the  fuse-wire. 

General  laws.  In  interpreting 
the  results  of  such  measurements 
and  in  calculating  the  heat  of  other 
reactions  from  data  collected  in 
this  way  two  general  theorems 
are  of  constant  application. 

1.  Heat  of  formation  of  a  substance  equal  to  its  heat  of  decompo- 
sition.   By  the  heat  of  formation  of  a  substance  is  meant  the  heat 
evolved  or  absorbed  in  the  formation  of  a  gram-molecular  weight  from 
its  elements.    Save  in  the.  case  of  rather  simple  substances  it  is  not 
possible  to  determine  the  heat  of  formation  directly.    By  decomposing 
a  substance  in  various  ways  and  determining  the  heat  changes  taking 
place,  it  is  often  possible  to  calculate  the  heat  of  formation  on  the 
assumption  that  it  is  equal  to  the  heat  of  decomposition.    This  is,  of 
course,  a  special  application  of  the  law  of  conservation  of  energy. 

2.  The  law  of  heat  summation;  the  law  of  Hess.    As  early  as  1840 
Hess  worked  out  the  general  theorem  that  the  heat  given  off  in  a  series 
of  transformations  is  independent  of  the  several  steps,  and  depends 
only  upon  the  initial  and  final  states  of  the  substance.    For  example, 
carbon  may  first  be  burned  to  carbon  monoxide,  and  this  in  turn  to 
the  dioxide.    The  sum  of  the  heats  in  the  two  stages  is  the  same  as  if 
the  carbon  were  to  be  burned  directly  to  the  dioxide.    This  principle 
is  of  constant  application  in  calculating  some  step  in  the  series  which 
cannot  be  directly  measured.    Thus  the  heat  of  formation  of  carbon 


FIG.  121 


336  GENERAL  CHEMISTRY 

monoxide  cannot  be  directly  determined,  but  may  be  calculated  from 
the  equations  C  +  O2  =  C02  +  96,960  cal. 

CO  +  O  =  CO2  +  67,960  cal. 

The  difference  between  these  two  equations  gives  us  the  following : 
C  +  O  =  CO  +  29,000  cal. 

This  law  is  also  a  special  application  of  the  law  of  conservation  of  energy, 
but  was  formulated  before  the  more  general  law  was  well  recognized. 

Types  of  heat  measurements.  The  general  scope  of  heat  measure- 
ments may  be  indicated  by  discussing  a  number  of  typical  cases. 

1.  Heat  of  combustion.  One  of  the  most  important  determinations 
is  that  of  the  heat  evolved  when  a  substance  is  burned  to  its  ultimate 
oxidation  products.  Such  measurements  are  of  the  greatest  technical 
importance  as  well  as  of  much  theoretical  value.  As  examples  we 
have  the  following: 

Methane    .  ,  .  CH4  +  2  O2  =  CO2  +  2  H2O  +  211,930  cal. 

Benzene     .  .  .  C6H6  +  1\  O2  =  6  CO2  +  3  H2O  +  801,160  cal. 

Alcohol  .  '.  .  .  C2H5OH  +  3  O2  =  2  CO2  +  3  H2O  +  341,790  cal. 

Water  gas  .  .  CO  +  H2  +  1J  O2  =  CO2  +  H2O  +  126,029  cal. 

Fuels.  The  various  materials  used  as  fuels  differ  much  in  the  heat 
which  they  give  out  when  burned.  While  many  other  factors  are 
concerned  in  the  value  of  a  fuel,  the  chief  one  is  its  heat  of  com- 
bustion. The  heat  evolved  by  the  combustion  of  one  gram  of  a  fuel 
is  called  its  calorific  value.  In  large  contracts  the  price  paid  for  a 
fuel  is  generally  based  on  its  calorific  value,  as  well  as  upon  its 
adaptability  to  the  use  to  which  it  is  to  be  put.  The  following  table 
will  give  some  average  values  for  a  few  common  fuels : 

CALORIFIC  VALUES   OF   FUELS 

Wood  (air-dried) •    ,  • about  3800-4000  cal 

Lignite  (brown),  8%  ash,  12%  moisture about  5400  cal. 

Bituminous  coal  (Pennsylvania),  35%  volatile  matter,  6%  ash     .  about  8300  cal. 

Bituminous  coal  (Pocahontas),  18%  volatile  matter,  6%  ash   .     .  about  8700  cal. 

Anthracite  coal  (Connellsville),  12%  ash about  7300  cal. 

Coke,  10%  ash about  7300  cal. 

Foods.  One  of  the  most  important  functions  of  food  is  to  supply  the 
energy  expended  by  the  body.  Apart  from  mineral  salts  and  water, 
most  of  the  constituents  of  foods  which  are  digested  ultimately  undergo 
oxidation  in  the  body,  the  carbon  and  hydrogen  being  in  large  part 
oxidized  into  carbon  dioxide  and  water.  The  heat  of  the  body  is  due 


THERMOCHEMISTRY  337 

to  this  oxidation.  The  chief  function  of  a  considerable  portion  of  our 
food  is  to  maintain  this  supply  of  energy,  so  the  calorific  value  of  foods 
is  a  matter  of  much  importance.  It  is  estimated  by  Sherman  that  a  man 
of  average  size,  living  a  normal  professional  life  involving  no  manual 
labor,  requires  the  supply  of  from  2,000,000  to  2,250,000  cal.  daily. 

On  the  average  the  calorific  value  of  the  three  principal  groups  of 
foodstuffs,  as  determined  in  the  calorimeter  and  in  actual  combustion 
within  the  body,  are  as  follows : 


CALORIMETER 

BODV  COMBUSTION 

Carbohydrates 

4100  cal. 

4000  cal. 

Fats  

9450  cal. 

9000  cal. 

Proteins 

5650  cal. 

4000  cal. 

The  following  table,  also  taken  from  Sherman,  shows  the  weight  in 
grams  of  a  few  important  foods  required  to  yield  100,000  cal. 

Beef,  free  from  visible  fat      .  86  g.          Butter 14  g. 

Beef,  round  steak     ....  64  g.          Bread  (white) 38  g. 

Bacon,  smoked 19  g.           Sugar 25  g. 

Eggs 67  g.          Potatoes 120  g. 

Milk 145  g.          Beans  (dried) 29  g. 

2.  Heat  of  formation.  Knowledge  as  to  the  heat  of  formation  of  a 
substance  is  often  of  the  greatest  service  in  forecasting  its  probable 
conduct  toward  various  reagents.  If  this  heat  is  great,  the  substance 
is  likely  to  be  indifferent  toward  decomposing  reagents,  or  to  act  upon 
them  only  at  high  temperatures.  Thus,  the  heat  of  formation  of  carbon 
dioxide  is  96,960  cal.  This  large  heat  value  would  indicate  that  only 
those  substances  which  have  a  very  strong  affinity  for  oxygen  will  be 
able  to  reduce  the  oxide. 

In  the  case  of  some  simple  substances  the  heat  of  formation  can  be 
determined  directly,  as  in  the  case  of  carbon  dioxide  and  of  sulfur. 
In  the  majority  of  cases,  however,  the  value  must  be  calculated  from 
measurements  on  the  heat  of  combustion.  For  example,  it  is  not 
possible  to  prepare  marsh  gas  from  its  elements  under  conditions 
which  permit  of  measuring  directly  its  heat  of  formation.  This  can 
be  calculated  from  the  following  equations : 

CH4  +  2  O2  =  CO2  +  2  H2O  +  211,930  cal.       (1) 

2  H2  +  02  =  2  H20  +  136,720  cal.  (2) 

C  +  O2  =  CO2  +  96,960  caL  (3) 


338  GENEKAL  CHEMISTRY 

By  adding  equations  (2)  and  (3)  we  get  233,680  cal.  as  the  heat 
evolved  in  the  formation  of  2  H2O  +  CO2  from  the  elements,  as 
against  211,930  cal.  when  formed  by  the  combustion  of  marsh  gas. 
Consequently,  233,680-211,930  cal.  =  21,750  cal.  must  be  evolved 
when  marsh  gas  is  formed  from  carbon  and  hydrogen. 

3.  Heat  of  solution.    When  a  substance  dissolves  in  water,  there  is 
always  a  change  in  the  temperature  of  the  water,  indicating  either 
a  heat  evolution  or  absorption  in  the  process  of  solution.    This  change 
is  the  algebraic  sum  of  a  number  of  separate  effects.   A  solid  or  a  gas 
after  solution  is,  in  a  sense,  in  the  state  of  a  liquid,  so  that  there  is  an 
effect  corresponding  in  a  general  way  to  the  heat  of  fusion  or  of  lique- 
faction.   There  is  usually  a  change  in  volume,  involving  mechanical 
work.    If  the  solute  is  an  electrolyte,  ionization  takes  place,  and  this, 
like  all  chemical  actions,  involves  heat  changes,  which  may  be  either 
positive  or  negative.    Undoubtedly  some  solutes  combine  chemically 
with  the  water,  forming  hydrates. 

It  is  not  possible  to  analyze  the  total  heat  change  into  the  fractions 
to  be  assigned  to  each  of  these  causes,  but  the  total  effect  is  of  much 
importance,  since  it  must  be  applied  as  a  correction  in  many  measure- 
ments in  which  the  reaction  takes  place  in  solution. 

4.  Reversible  reactions.  In  Chapter  XVI  it  was  stated  that  when  a 
reaction  is  taking  place  at  a  certain  temperature,  a  rise  of  10°  usually 
about  doubles  its  velocity.    A  reaction  which  evolves  heat  should 
therefore  proceed,  with  constantly  increasing  velocity,  to  completion. 
When  it  fails  to  do  so,  and  results  in  an  equilibrium,  it  is  evident  that 
there  must  be  some  simultaneous  occurrence  which  absorbs  the  heat 
of  reaction,  and  we  should  at  once  suspect  that  this  is  the  reverse 
reaction,  which  is  endothermic.    This  is  found  to  be  the  case,  one  reac- 
tion in  an  equilibrium  being  exothermic  and  the  other  endothermic. 
Equilibrium  will  result  when  the  heat  change  in  the  one  reaction  is 
just  balanced  by  that  in  the  other. 

These  statements  may  be  easily  tested  by  two  experiments  suggested  by  the 
equilibrium  equation 

NaCl  +  H2SO4  — >•  NaHSO4  +  HC1 

When  sodium  chloride  is  treated  with  moderately  concentrated  sulfuric  acid,  there 
is  at  the  outset  a  very  considerable  evolution  of  heat,  as  shown  by  a  thermometer 
introduced  into  the  solution.  On  the  other  hand,  when  sodium  hydrogen  sulfate 
is  treated  with  concentrated  hydrochloric  acid,  there  is  a  corresponding  cooling 
of  the  solution,  which  is  often  sufficient  to  cause  the  beaker  containing  the  solu- 
tion to  freeze  fast  to  a  wet  block  of  wood  on  which  it  is  placed. 


THEKMOCHEMISTKY  339 

Limits  of  temperature  in  combustion.  When  it  is  desired  to  secure 
a  very  high  temperature  by  a  chemical  reaction,  as  by  the  combustion 
of  hydrogen  or  carbon,  the  most  favorable  conditions  are  that  the 
reaction  shall  take  place  as  rapidly  as  possible  and  under  such  cir- 
cumstances as  will  involve  the  least  loss  of  heat  through  radiation 
and  conduction.  In  other  words,  the  greatest  quantity  of  heat  must 
be  liberated  in  the  least  possible  time  and  in  the  presence  of  as  little 
conducting  material  as  possible.  In  this  way  the  combustion  of  hydro- 
gen will  produce  a  temperature  sufficient  to  melt  platinum  and  even 
to  reach  2400°.  From  the  known  heat  of  reaction 
2  H2  +  02  =  2  H20  +  116,138  cal. 

it  would  seem  that  a  higher  temperature  should  be  secured,  for  in 
the  reaction  expressed  by  the  equation 

4  Al  +  3  O2  =  2  A12O3  +.  760,400  cal. 

a  temperature  estimated  at  as  much  as  4000°  is  obtainable.  The  dif- 
ference in  the  two  cases  lies  in  the  fact  that  the  one  reaction  is  revers- 
ible and  the  other  is  not.  In  the  combustion  of  hydrogen,  as  soon 
as  a  temperature  of  about  1200°  is  reached,  the  steam  formed  begins 
to  decompose,  as  indicated  in  the  equation 


The  higher  the  temperature  the  larger  the  percentage  of  decomposi- 
tion (see  table,  p.  61),  and  when  this  becomes  considerable,  the 
absorption  of  heat  prevents  any  further  rise  in  temperature.  In  the 
combustion  of  carbon  the  case  is  similar.  The  carbon  dioxide  decom- 
poses at  a  higher  temperature,  as  shown  in  the  equation 

2  CO2  ^=^  2  CO  +  O2 

This  reverse  reaction  absorbs  the  heat  of  the  direct  formation.  With 
aluminium  the  product  of  combustion  (A12O3)  is  stable  up  to  the 
extreme  limit  of  experiment,  and  all  of  the  heat  of  combination  is 
available  to  raise  the  temperature  of  the  resulting  aluminium  oxide 
and  surrounding  bodies. 

As  a  determining  factor  in  the  temperature  which  can  be  secured 
by  a  given  reaction  it  is  therefore  not  only  the  heat  of  reaction  which 
comes  into  account,  but  also  the  stability  of  the  reaction  product  at 
high  temperatures.  The  setting  in  of  a  reverse  reaction  always  puts 
a  limit  upon  the  temperature  which  can  be  produced  by  the  reaction. 


340  GENERAL  CHEMISTRY 

Conditions  for  stability  of  exothermic  and  endothermic  compounds. 
The  dissociation  of  the  strongly  exothermic  compounds,  water  and 
carbon  dioxide,  at  high  temperatures  is  in  accord  with  the  principle 
of  Le  Chatelier  (p.  202),  which  predicts  that  a  heat-absorbing  reac- 
tion -will  tend  to  take  place  as  the  temperature  is  raised.  It  has  been 
found  to  be  true  in  general  that  exothermic  compounds  tend  to  dissoci- 
ate in  a  similar  way  at  high  temperatures,  and,  conversely,  that  they 
become  more  stable  as  the  temperature  is  lowered.  In  view  of  the  large 
heat  evolution  attending  the  formation  of  aluminium  oxide,  we  should 
expect  it  to  dissociate  when  heated  sufficiently,  and  it  doubtless  would 
do  so  if  we  could  reach  the  necessary  temperature. 

On  the  other  hand,  the  principle  of  Le  Chatelier  would  lead  us  to 
expect  that  an  endothermic  compound  would  tend  to  form  from  its 
elements  at  a  high  temperature,  but  would  be  unstable  at  lower  tem- 
peratures, dissociating  with  heat  evolution.  This  is  found  to  be  the  case. 
For  example,  carbon  and  hydrogen  readily  combine  at  the  temperature 
of  the  electric  arc  to  form  acetylene,  and  this  compound  becomes  more 
and  more  unstable  as  the  temperature  is  lowered.  It  is  easy  to  see 
that  if  this  world  was  once  in  gaseous  condition,  the  first  compounds 
to  form  must  have  been  endothermic  in  character. 


CHAPTER  XXV 

THE  SILICON  FAMILY  AND  BORON 

Introduction.  In  both  families  of  Group  IV  of  the  periodic  arrange- 
ment, the  elements  of  small  atomic  weight  are  acid-forming  in  charac- 
ter and  resemble  carbon  in  many  particulars,  but  with  increase  in 
atomic  weight  they  become  metallic  in  their  properties.  It  is  therefore 
convenient  to  describe  the  acid-forming  elements,  silicon,  titanium,  and 
zirconium,  at  this  point,  reserving  the  more  metallic  elements  for  a 
subsequent  chapter.  It  will  be  appropriate  to  describe  the  element 
boron  in  connection  with  silicon,  since  it  also  has  acid-forming  prop- 
erties, while  the  other  members  of  the  third  group  are  metallic  in 
character. 

SILICON 

Occurrence.  Next  to  oxygen,  silicon  is  the  most  abundant  element. 
Neither  the  element  nor  its  compounds  are  found  in  the  air,  nor  to 
any  considerable  extent  in  water,  but  the  solid  crust  of  the  earth  is 
estimated  to  contain  28  per  cent  of  silicon.  All  varieties  of  granite, 
gneiss,  sandstone,  shale,  clay,  and  marl  contain  large  percentages 
of  the  element,  limestone  and  dolomite  being  the  only  important 
geological  formations  measurably  free  from  it.  In  the  realm  of  inor- 
ganic nature  it  is  the  central  element  just  as  carbon  is  of  fundamental 
importance  in  organic  nature.  To  some  extent  its  compounds  are 
assimilated  by  plants  and  animals,  and  they  constitute  the  outer 
shell  of  many  aquatic  organisms. 

The  element.  In  the  laboratory,  crystallized  silicon  is  best  prepared 
by  the  reduction  of  the  dioxide  with  aluminium  powder: 
3  SiO2  +  4  Al  =  3  Si  +  2  A12O8 

The  silicon  dissolves  in  the  excess  of  melted  aluminium,  and  when  the 
solution  has  cooled  and  become  solid,  the  aluminium  is  dissolved  in 
hydrochloric  acid,  the  silicon  being  left  in  the  form  of  shining  metal- 
lic needles.  The  reduction  of  the  dioxide  with  carbon  has  always  pre- 
sented* the  difficulty  that  the  reduced  element  tends  to  combine  with 
excess  of  carbon  to  form  a  carbide.  This  difficulty  has  been  overcome 

341 


342  GENERAL  CHEMISTRY 

to  a  large  extent,  and  nearly  pure  silicon  is  now  manufactured  in 
large  quantities.  By  reducing  a  mixture  of  the  oxides  of  silicon  and- 
iron with  carbon,  an  alloy  of  the  two  elements  called  ferrosilicon  is 
obtained.  This  alloy,  as  well  as  the  purer  silicon,  finds  an  important 
application  in  the  metallurgy  of  iron. 

Properties.  The  element  presents  a  close  analogy  with  carbon  in 
that  it  can  be  obtained  in  amorphous  form,  as  well  as  in  crystals  of 
the  isometric  system,  resembling  diamond.  The  crystals  are  very  hard, 
easily  scratching  glass,  and  have  a  density  of  2.35.  They  melt  at 
about  1450°.  A  lump  of  the  element  is  very  brittle  and  breaks  with 
a  crystalline  fracture,  which  has  a  metallic,  silvery  appearance. 

Chemical  conduct.  Silicon  is  readily  attacked  by  the  halogens,  form- 
ing gaseous  compounds.  The  gaseous  halogen  hydrides  attack  it  at 
red  heat,  with  evolution  of  hydrogen.  The  alkalies  dissolve  it,  evolv- 
ing hydrogen  and  forming  soluble  silicates : 

4  NaOH  +  Si  =  Na4SiO4  -f  2  H2 

Water  containing  a  trace  of  alkalies  acts  slowly  upon  it  in  a  similar 
way.  It  is  not  attacked  by  oxygen  as  easily  as  would  be  expected, 
considering  the  difficulty  with  which  its  oxide  is  reduced.  This  in- 
activity is  due  to  its  high  melting  point  and  to  the  fact  that  a  thin 
film  of  nonvolatile  oxide  forms  upon  the  surface,  which  protects  it 
from  further  action.  At  high  temperatures  it  combines  with  the  great 
majority  of  elements,  forming  silicides  such  as  those  of  magnesium 
(Mg2Si)  and  carbon  (CSi). 

Halogen  and  hydrogen  compounds  of  silicon.  A  large  number  of 
compounds  of  silicon  with  hydrogen  and  the  halogens  have  been  pre- 
pared, and  they  are  of  interest  as  showing  the  close  relationship  of  the 
element  with  carbon.  The  hydrides  are  formed  by  the  action  of  acids 
or  water  upon  a  suitable  silicide,  just  as  are  those  of  carbon  (C2H2)  and 
sulfur  (H2S)  :  Mg^  +  4  H2O  =  SiH4  +  2  Mg(OH)2 

The  halogen  derivatives  are  formed  by  direct  union  of  silicon  with 
the  halogens,  or  by  conducting  the  halogen  vapor  through  a  heated 
mixture  of  carbon  and  silicon  oxide: 

Si02  +  2  C  +  2  C12  =  SiCl4  +  2  CO 

The  following  table  will  serve  to  show  the  properties  of  a  number 
of  these  silicon  compounds  and  those  of  the  corresponding  carbon 
compound. 


THE   SILICON  FAMILY  AND  BORON 


343 


TABLE  OF  SOME  CARBON  AND  SILICON  COMPOUNDS 


FORMULA 

MELTING 

POINT          « 

BOILING 
POINT 

FORMULA 

MELTING 
POINT 

BOILING 
POINT 

CH4.     .     . 

-184. 

-164. 

SiH4  .     .     . 

-200. 

C2H6      .     . 

-  171.4 

-85.4 

Si2H6     .     . 

-138. 

52. 

CC14  ... 

-23.8 

76.7 

SiCl4      .     . 

-102. 

57.5 

C2C16      .     . 

187. 

185. 

Si2Cl6     .     . 

-1. 

146. 

CSC18      .     . 

160. 

268. 

Si3Cl8     .     . 

-12. 

210. 

CBr4  .     .     . 

92.5 

189.5 

SiBr4      .     . 

-12. 

153. 

CI4    .     .     . 

decomp. 

SiI4   .     .     . 

120.5 

290.  ± 

CF4   .     .     . 

-  15. 

SiF4  .     .     . 

-102. 

sublimes 

CHF3     .     . 

SiHF3    .     . 

-110. 

30.2 

CHC13    .     . 

-60.3 

61.3 

SiHCl3  .     . 

-134. 

33. 

CHBr3   .     . 

9. 

146. 

SiHBr3  .     . 

-60. 

115. 

CHI3      .     . 

119. 

decomp. 

SiHI3     .     . 

8. 

decomp. 

Silicon  fluoride  (SiFJ.  One  of  these  compounds,  silicon  tetrafluoride 
(SiF4),  deserves  special  mention.  It  is  a  gas  at  ordinary  temperature, 
and  is  easily  prepared  by  the  action  of  hydrofluoric  acid  upon  silicon 
dioxide  :  2  HQ 


In  a  somewhat  similar  way  hydrofluoric  acid  acts  upon  silicates,  con- 
verting both  the  silicon  and  the  metals  into  fluorides  (p.  247). 

Fluosilicic  acid  (H2SiF6).  When  silicon  fluoride  is  conducted  into 
water,  it  is  decomposed  according  to  the  following  equation  : 

SiF4  +  4  H20  =  Si(OH)4  +  2  H2F2 

A  part  of  the  silicon  hydroxide  then  dissolves  in  the  hydrofluoric  acid  : 
Si(OH)4  +  3  H2F2  =  H2SiF6  +  4  H2O 

The  soluble  product  H2SiFg  is  called  fluosilicic  acid.  It  is  a  moder- 
ately strong  acid  and  forms  soluble,  well-crystallized  salts  with  most 
of  the  metals.  The  potassium  salt  K2SiFg  is  very  sparingly  soluble, 
so  that  the  acid  is  frequently  employed  as  a  test  for  potassium.  Most 
of  those  elements  which  in  the  elementary  state  resemble  metals,  but 
whose  chemistry  is  that  of  nonmetals,  form  fluo-acids  of  the  same  gen- 
eral kind.  Among  these  are  all  the  elements  in  this  group  except  carbon. 

Preparation  of  fluosilicic  acid.  A  dilute  solution  of  fluosilicic  acid  is  conveniently 
prepared  by  the  use  of  the  apparatus  shown  in  Fig.  122.  A  mixture  of  sand 
(SiO2)  and  powdered  fluor  spar  is  placed  in  the  flask  A,  and  concentrated  sulf  uric 
acid  is  added.  The  hydrofluoric  acid  liberated  acts  upon  the  sand,  forming  silicon 
tetrafluoride.  This  is  conducted  through  a  wide  delivery  tube  (jB)  into  a  beaker 


344 


GENERAL   CHEMISTRY 


FIG.  122 


of  water,  provided  with  a  layer  of  mercury  on  the  bottom,  into  which  the  delivery 
tube  dips.  This  is  to  prevent  the  solid  silicic  acid  which  also  forms  from  clog- 
ging the  tube.  At  the  close  of  the 
operation  the  silicic  acid  is  filtered  off. 
The  filtrate  contains  the  fluosilicic 
acid  in  solution. 

Silicon  dioxide  (silica)  (Si02). 

Although  several  oxides  of  sili- 
con have  been  described,  the 
dioxide  SiO2,  called  silica,  is  the 
only  one  which  is  at  all  well  char- 
acterized. Practically  all  of  the 
silicon  of  nature  occurs  either  as 
the  dioxide  itself  or  as  its  deriva- 
tives, so  that  it  is  a  most  impor- 
tant substance.  As  found  in  nature,  silica  is  usually  in  the  crystalline 
variety  known  as  quartz.  It  forms  beautiful  colorless  crystals  belong- 
ing to  the  hexagonal  system,  which  are  sometimes  of  great  size.  A 
single  one  found  in  California  weighed  over  a  ton.  These  crystals  are 
frequently  tinted  by  impurities,  producing  smoky  and  milky  quartz  and 
amethyst.  An  examination  of  well-formed  crystals  shows  that  they  are 
not  symmetrical  but  are  rights  and  lefts  like  a  pair 
of  gloves.  Fig.  123  shows  such  a  pair.  Above  800° 
quartz  changes  into  two  other  forms,  known  as  tridy- 
mite  and  cristobalite,  and  there  are  several  other 
crystal  forms  known.  Quartz  is  a  constituent  of  many 
conglomerate  rocks,  such  as  sandstone.  Silica  also 
occurs  in  amorphous  form,  usually  impure  and  often  partially  hydrated. 
Among  such  forms  are  flint  and  opal.  It  also  acts  as  the  binding  mate- 
rial which  unites  the  several  minerals  constituting  granite  and  gneiss. 
Quartz  has  a  density  of  2.66  and  is  hard  enough  to  scratch  glass 
and  most  metals.  It  is  therefore  used  in  the  form  of  grindstones  and 
powder  for  grinding  and  polishing  purposes.  Silica  is  very  difficult  to 
melt,  but  at  the  temperature  of  the  oxyhydrogen  blowpipe  it  softens 
to  a  viscous  liquid  resembling  melted  glass,  which  can  be  drawn  into 
threads  or  fashioned  into  various  laboratory  utensils,  such  as  crucibles 
or  flasks,  which  have  many  desirable  qualities.  They  are  not  attacked 
by  most  reagents,  and,  owing  to  the  fact  that  silica  has  a  very  small 
coefficient  of  expansion  with  temperature,  such  vessels  can  be  heated 
to  redness  and  plunged  into  water  without  danger  of  cracking. 


FIG.  123 


THE  SILICON  FAMILY  AND  BOKON  345 

The  melting  point  of  crystallized  silica.  Silica  is  undoubtedly  a  very 
highly  crystalline  substance,  yet  it  would  appear  to  have  no  sharp 
melting  point  but  to  pass  by  insensible  stages  into  a  viscous  liquid. 
On  cooling,  this  gradually  hardens  without  resuming  a  crystalline 
structure.  This  conduct  does  not  seem  to  be  in  accord  with  the  gen- 
eral principles  already  developed  concerning  the  melting  point  of  crys- 
tals (see  p.  79),  but  the  explanation  is  really  very  simple.  The  melted 
silica  is  so  viscous  that  it  retains  the  form  of  the  original  solid,  and 
this  amorphous  liquid,  on  further  heating,  gradually  softens.  On  cool- 
ing, the  viscous  liquid  permanently  supercools  and,  as  a  rule,  fails  to 
crystallize  again.  By  optical  observations  it  has  been  shown  that  the 
true  melting  point  is  about  1600°. 

Acids  of  silicon.  Like  carbon  dioxide,  silica  is  an  acid  anhydride, 
and  this,  together  witn  its  action  with  hydrofluoric  acid,  constitutes 
its  most  interesting  chemical  property.  In  combination  with  water  it 
forms  not  only  one,  but  a  considerable  variety  of  acids. 

Orthosilicic  acid  (H4Si04).  Since  silicon  is  tetravalent,  its  normal 
hydroxide  would  have  the  formula  Si(OH)4.  This  appears  to  be  a 
tetrabasic  acid,  for  salts  such  as  Ca2SiO4  and  KAlSiO4  are  well  known. 
This  hydroxide  is  called  orthosilicic  acid,  the  formula  being  written 
H4SiO4,  and  its  salts  are  called  orthosilicates. 

Metasilicic  acid  (H2Si03).  Orthosilicic  acid  readily  loses  one  mole- 
cule  of  water:  H4SiO4  =  H£iO,  +  H,O 

The  acid  so  formed  is  analogous  to  carbonic  acid  (H2CO3)  and  is  called 
metasilicic  acid.  Its  salts  are  called  metasilicates,  the  sodium  and  cal- 
cium salts  having  the  formulas  Na2SiO3  and  CaSiO3. 

Polysilicic  acids.  There  is  another  way  in  which  the  orthosilicic 
acid  may  lose  water,  namely,  by  the  cooperation  of  more  than  one 
molecule.  This  may  be  understood  by  reference  to  the  following 
structural  formulas: 


o 

Si^OH 
O 


Si^OH  SitS 


H6Si207 — +    H4Si206 — *     H2Si2O6 — *     2  SiO? 


346  GENEKAL  CHEMISTRY 

With  three  molecules  we  might  have,  in  a  similar  way, 
3  Si(OH)4— >H8Si8010— *H6Si809— ^H,Si808^H2Si30,^3  SiO2 

When  two  or  more  molecules  of  an  acid  become  condensed  in  this 
way,  forming  products  such  as  the  ones  represented  above,  the  result- 
ing acids  are  called  condensed  acids,  or  polyacids,  the  ones  just  formu- 
lated being  polysilicic  acids.  To  a  greater  or  less  extent  this  tendency 
is  observed  with  nearly  all  oxygen  acids,  and  we  have  already  had  an 
illustration  in  pyrosulfuric  acid  (H2S2O7). 

In  the  case  of  silicon  this  tendency  is  very  pronounced,  and  most 
of  the  natural  silicates  are  salts  of  such  acids.  While  there  is  no 
doubt  as  to  the  existence  of  these  various  acids,  none  of  them,  not 
even  the  ortho  or  meta  acids,  have  been  prepared  in  pure  form.  They 
all  gradually  lose  water  and  pass  finally  into  the  dioxide. 

The  silicates.  Almost  all  the  silicates,  whether  salts  of  simple  or 
of  condensed  acids,  are  insoluble  in  water,  excepting  those  of  sodium 
and  potassium.  The  latter  may  be  prepared  by  fusing  pure  silica 
with  sodium  hydroxide  or  carbonate : 

4  NaOH  +  SiO2  =  NaJ3iO4  +  2  H2O 
Na2C03  +  Si02  =  Na2Si03  +  CO2 

A  solution  of  the  products  of  such  fusions  in  water  is  called  water 
glass.  It  doubtless  contains  a  mixture  of  the  various  silicates  of 
sodium  or  potassium.  When  the  solution  is  evaporated,  it  forms  a 
thick  liquid,  which  gradually  hardens  into  a  glass.  It  is  used  as  a 
cement  or  glue  for  many  purposes. 

When  salts  of  the  various  metals  are  added  to  a  solution  of  sodium 
silicate,  the  insoluble  silicates  are  precipitated : 

Na2SiO3  +  CaCl2  =  CaSiO3  +  2  NaCl 

When  most  natural  silicates  are  fused  with  sodium  carbonate,  they 
are  decomposed,  and  sodium  silicate  is  formed : 

CaSi03  +  Na2C08  =  CaCO3  +  Na2SiO8 

The  melted  sodium  carbonate  forms  a  liquid  in  which  the  silicates  are  soluble 
but  the  carbonates  are  not.  The  insoluble  carbonate  is  therefore  precipitated. 
When  the  melt  is  cooled  and  digested  with  water,  the  sodium  silicate  dissolves, 
while  the  carbonate  and  oxides  of  the  other  metals  are  left  undissolved.  They 
may  be  filtered  off  and  dissolved  in  acids.  Such  fusions  are  very  frequently 
resorted  to  in  chemical  analysis  as  a  means  of  decomposing  the  silicates  and  get- 
ting their  constituents  into  solution,  as  many  of  them  are  not  attacked  by  acids. 


THE  SILICON  FAMILY  AND  BORON 


347 


Varieties  of  natural  silicates.  Some  natural  silicates  are  salts  of 
metasilicic  acid ;  others  are  derived  from  orthosilicic  acid.  More  fre- 
quently they  are  salts  of  various  polysilicic  acids.  These  are  designated 
by  prefixes,  such  as  di-  and  tri-,  according  to  the  number  of  molecules 
of  the  orthosilicic  acid  which  have  been  condensed.  The  disilicic  acid 
H4Si2O6  and  the  trisilicic  acid  H4SigO8  are  the  most  common  of  these. 
Very  often,  when  the  acid  has  as  many  as  four  hydrogen  atoms, 
two  or  more  different  metals  replace  the  hydrogen,  forming  mixed 
salts,  such  as  KAlSiO4  (mica)  or  KAlSi3O8  (feldspar).  The  following 
table  will  illustrate  these  types  of  silicates : 


VARIETIES 

ACIDS 

SALTS 

MINERALS 

Metasilicates 

H2Si03 

CaSiO3 

Wollastonite 

Orthosilicates 

KLSiO, 

/Mg2Si04 

Olivine 

\Zn2SiO4 

Willemite 

Disilicates 

JH6Si207 

Mg3Si207 

Serpentine 

|H4Si206 

KAlSi206 

Leucite 

Trisilicates 

H4Si308 

fKAlSi3O8 

Orthoclase 

LNaAlSi308 

Albite 

Fusion  of  the  silicates.  Some  of  the  silicates  have  a  sharp  melting 
point  and  recrystallize  on  cooling.  Many  of  those  with  a  low  melting 
point  resemble  silica  in  that  they  fuse  into  a  viscous  liquid  which,  on 
cooling,  hardens  without  crystallizing.  If  several  different  silicates  are 
melted  together,  they  mix  freely  to  a  homogeneous  liquid.  Crystals  of 
definite  composition  may  separate  from  the  melt  when  it  is  cooled, 
but  more  frequently  the  liquid  solution  simply  becomes  more  and 
more  viscous,  until  it  is  as  rigid  as  a  true  solid.  Such  products  are 
called  glasses,  and  they  are  to  be  regarded  as  very  viscous  solutions 
of  one  silicate  in  another.  The  industry  of  glass  making  will  be 
referred  to  in  a  subsequent  chapter  (p.  453). 

Colloidal  silicic  acid.  When  a  rather  concentrated  solution  of  water 
glass  is  treated  with  an  excess  of  strong  acid,  silicic  acid  is  liberated  as 

a  jellylite  maSS  =  Na4Si04  +  4  HCl  =  4  Nad  +  H4Si04 . 

In  dilute  solution  there  is  no  apparent  change  on  adding  the  acid,  and 
it  would  be  inferred  that  silicic  acid  is  moderately  soluble  in  water. 
If,  however,  the  acidulated  solution  is  placed  in  a  vessel  A,  the 
bottom  of  which  is  formed  of  parchment  (like  a  tambourine),  and  this 
vessel  is  in  turn  placed  in  a  larger  one  (J5),  containing  water,  as  shown 


348 


GENERAL   CHEMISTRY 


FIG.  124 


in  Fig.  124,  the  excess  of  hydrochloric  acid  and  the  sodium  chloride 
formed  in  the  reaction  pass  through  the  parchment,  but  the  silicic 
acid  does  not.  If  the  water  in  the  outer  vessel  is  renewed  from  time 
to  time,  tHe  silicic  acid  alone  is  left  in  solution  in  the  inner  vessel. 

Such  an  arrangement  is  called  'a  dia- 

lyzer,  and  the  process  is  called  dialysis. 
The  solution  of   silicic  acid  looks 

perfectly  clear,  but  that  it  is  not  a 
—     true   solution    is  shown  by   the  fact 

that  its  freezing  point  is  practically 
the  same  as  that  of  pure  water.  A  beam  of  light  from  a  projecting  lan- 
tern passed  through  the  liquid  makes  a  bright  path,  as  it  does  in  a 
dusty  room  (Fig.  125),  whereas  it  does  not  do  this  with  an  ordinary 
solution.  The  silicic  acid  cannot  be  filtered  out  of  such  a  liquid  by 
ordinary  filter  paper,  yet  is  present  as  a  kind  of  suspension.  All  this 
conduct  is  characteristic  of  a  colloid  (p.  132),  and  the  acids  of  silica 
exhibit  it  to  a  high  degree.  As  long  as  no  silicic  acid  separates,  the 
liquid  is  called  the  hydrosol.  If  we  attempt  to  concentrate  it  past  a  cer- 
tain point,  it  separates  as  a  jelly,  called  the  Jiydrogel.  The  hydrogel  may 
also  be  separated  by  the  addition  of  al- 
most any  electrolyte,  like  common  salt. 
Colloidal  silicic  acids  are  present  in 
many  natural  waters,  being  derived 
from  the  hydrolysis  of  silicates  or  from 
the  action  of  acids  upon  them.  In  some 
localities,  as  in  Yellowstone  Park,  such 
waters  gradually  deposit  the  hydrogel, 
which  then  loses  water  and  builds  up  a 
basin  of  silica  about  the  spring.  Organic 
materials  immersed  in  such  waters  have 

silica  deposited  all  through  their  structure  and  become  petrified.  In 
Arizona  fallen  forests  of  petrified  logs  are  found  turned  into  agate  in 
this  way,  yet  retaining  the  structure  of  the  original  wood. 

Silicon  carbide  (carborundum)  (SiC)  .  When  sand,  coke,  and  a  little 
sodium  chloride  are  heated  in  an  electric  furnace,  the  silica  is  reduced, 
the  resulting  silicon  combining  with  carbon  to  form  a  carbide  : 


pIG  125 


Si  +  C  =  SiC 


THE   SILICON  FAMILY  AND  BORON 


349 


FIG.  126 


FIG.  127 


After  the  reaction  the  carbide  is  found  in  the  form  of  beautiful  crystal 
plates,  which  are  colorless  when  pure,  but  are  usually  a  brilliant  iri- 
descent purple,  the  color  being  largely  due  to  optical  rather  than  to 
chemical  causes.  The  material  is  almost  as  hard  as  diamond  and  is 
much  used  in  place  of  emery  for  grindstones  and  polishing  powders. 
Although  its  heat  of  combustion  is  great,  amounting  to  238,800  cal., 
it  is  extremely  inactive  chemically  except  at  very  high  temperatures. 
It  then  acts  as  a  strong  reducing  agent,  both  of  the  constituent  ele- 
ments being  converted  into  oxides. 

Commercial  production  of  carborundum.  The  furnace  in  which  carborundum  is 
prepared  is  quite  similar  to  the  one  employed  in  the  manufacture  of  graphite 
(p.  277).  It  is  built  out  of  loose  brick  and  half  filled  with  the  charge.  The 
electrodes  are  put  in  place, 
a  layer  of  pieces  of  carbon 
spread  between  them,  and 
the  charge  filled  in  to  the 
top.  A  cross  section  of  a 
furnace  so  charged  is  rep- 
resented in  Fig.  126.  A  is 
the  core  of  carbon,  B  the 
mixture  of  materials.  The 
broken  carbon  offers  great 
resistance  to  the  current,  and  the  heat  along  the  carbon  core  becomes  very  intense. 
This  causes  reaction  to  take  place  from  the  center  for  some  distance  into  the 
mass,  the  result  being  represented  in  Fig.  127.  A  is  the  core  of  carbon  sur- 
rounded by  a  core  of  crystallized  carborundum  B,  and  this  in  turn  by  a  layer 
(C)  of  amorphous  carborundum  and  partially  reduced  products.  These  are  known 
as  "  white  stuff,"  siloxicon,  or  silundum,  according  to  their  character.  They  are 
useful  for  making  furnace  linings  and  fire  brick.  The  material  D  is  unchanged. 

The  silicate  industries.  A  number  of  most  important  industries  are 
based  upon  the  silicates  and  the  chemical  transformation  which  they 
undergo.  The  oldest  of  these  are  the  various  ceramic  industries,  in- 
cluding the  making  of  bricks,  tile,  terra  cotta,  dishes,  porcelain,  and 
glass.  Glass  making,  with  its  many  details,  is,  from  a  chemical  stand- 
point, very  closely  associated  with  the  making  of  porcelain  and  the 
glazing  of  clay  vessels.  Cement  making  also  involves  similar  chemical 
principles.  These  three  great  industries  are  sometimes  collectively 
spoken  of  as  the  ceramic  industries,  and  their  discussion  would  natu- 
rally take  place  in  the  present  chapter,  but  since  all  these  industries 
are  based  on  the  conduct  of  metallic  silicates,  it  seems  best  to  delay 
the  discussion  until  a  number  of  the  metals  have  been  studied,  espe- 
cially sodium,  calcium,  and  aluminium. 


350  GENERAL  CHEMISTRY 

TITANIUM 

Occurrence.  Titanium  occurs  rather  sparingly  in  nature  and  is 
usually  found  in  the  form  of  the  dioxide  TiO2,  called  rutile,  or  as 
an  iron  titanite  (FeTiO3)  known  as  ilmenite,  or  as  a  variable  con- 
stituent of  certain  magnetic  iron  ores.  In  small  quantities  titanium 
is  very  widely  distributed  in  nature,  being  a  common  constituent  of 
many  minerals  and  being  found  in  traces  in  both  vegetable  and  animal 
organisms. 

The  element.  The  element  can  be  obtained  by  the  reduction  of  the 
dioxide  with  carbon  in  an  electric  furnace,  but  prepared  in  this  way 
it  always  contains  carbon  and  usually  nitrogen.  Very  pure  specimens 
have  been  prepared  by  the  action  of  titanium  chloride  on  sodium  in  a 
closed  steel  bomb :  TiC1^  +  4  Na  =  Ti  +  4  NaC1 

9 
When  the  element  contains  carbon,  it  is  hard  and  very  brittle.   When 

pure,  it  is  brittle  when  cold,  but  can  be  forged  at  a  low  red  heat. 
Its  melting  point  is  above  that  of  platinum,  being  about  1800°. 
Its  specific  gravity  is  4.5. 

The  compounds.  The  dioxide  of  titanium,  like  that  of  silicon,  is  an 
acid  anhydride  and  forms  a  large  number  of  acids  closely  resembling 
the  various  types  of  silicic  acids.  These  are  even  weaker  than  those 
of  silicon,  and  their  salts  hydrolyze  more  readily  and  form  colloidal 
solutions.  Fluotitanic  acid  (H2TiF6)  and  its  salts  are  well  known. 

Unlike  silicon,  titanium  also  forms  salts  in  which  the  element  acts 
as  a  trivalent  metal.  The  titanic  salts  are  formed  by  the  action  of 
nascent  hydrogen  upon  derivatives  of  the  dioxide.  The  sulfate 
Ti2(SO4)3  and  the  chloride  TiCl3  •  6  H2O  are  examples.  These  salts 
are  either  green  or  violet  in  color. 

At  high  temperatures  titanium  shows  a  very  marked  tendency  to 
unite  with  nitrogen,  the  nitride  TiN  being  the  product  of  this  direct 
union.  The  nitride  is  therefore  always  produced  in  any  attempt  to 
prepare  titanium  in  an  apparatus  to  which  air  has  access,  and  this 
compound  was  formerly  considered  to  be  the  element  itself.  When 
iron  ores  containing  titanium  are  smelted,  a  substance  resembling 
crystallized  copper  is  often  found  in  the  slag  or  adhering  to  the  lining 
of  the  furnace.  This  was  also  at  one  time  supposed  to  be  the  metal, 
but  is  now  known  to  have  the  formula  TiloC2Ng. 


THE  SILICON  FAMILY  AND  BOKON  351 

ZIRCONIUM 

Zirconium.  This  is  a  rather  rare  element  which  occurs  in  nature 
chiefly  as  the  compound  ZrSiO4,  called  zircon.  It  is  beautifully  crys- 
tallized in  octahedra  and  is  found  in  the  Carolinas,  Ceylon,  and  other 
localities,  especially  in  gold  and  diamond  sands.  Clear  specimens, 
usually  tinted  some  shade  of  yellow  or  red,  are  used  as  semiprecious 
stones,  under  the  name  hyacinth.  The  oxide  of  zirconium  (ZrO2)  is 
found  in  considerable  deposits  in  Brazil. 

While  zirconium  forms  acids  similar  to  those  of  silicon  and  titanium, 
particularly  the  fluozirconic  acid  H2ZrFg,  its  metallic  properties  are 
much  more  developed  and  it  forms  many  salts  in  which  the  element 
plays  the  part  of  a  tetravalent  metal  like  tin.  As  an  element  it  is 
therefore  intermediate  between  the  acid-forming  elements  of  this 
group  and  those  which  are  more  distinctly  metals. 

BORON 

The  trivalent  element  boron  finds  a  place  as  the  first  member  of 
Group  III  of  the  periodic  classification,  but  in  its  properties,  as  well 
as  in  those  of  its  compounds,  it  is  much  more  closely  related  to  sili- 
con than  to  the  other  trivalent  elements.  Since -it  is  trivalent,  the 
formulas  of  its  compounds  differ  from  those  of  silicon. 

Occurrence.  The  hydroxide  of  boron  (B(OH)3),  known  as  boric  acid, 
occurs  in  many  hot  springs,  particularly  in  Italy  and  California,  and 
this  was  formerly  the  chief  source  of  the  compounds  of  boron.  The 
element  is  also  found  in  large  quantities  as  borax  (Na2B4O?)  in  the 
desert  regions  of  California  and  Nevada.  Several  other  more  complex 
minerals  are  now  important  sources  of  boron  compounds.  Among 
these  are  a  complex  magnesium  salt  called  boracite,  found  in  the 
Stassfurt  deposits,  and  colemanite  (Ca2B6On  •  5  H2O),  found  in  a 
number  of  large  deposits  in  California. 

The  element.  Boron  has  an  extraordinary  affinity  for  oxygen,  so 
that  the  oxide  is  very  difficult  of  reduction.  By  heating  it  with  mag- 
nesium powder  an  impure  product  may  be  obtained  which  always 
contains  either  magnesium  or  oxygen : 

B2O8  +  3  Mg  =  2  B  +  3  MgO 

The  pure  element  has  been  obtained  by  Weintraub  by  the  action  of 
hydrogen  upon  boron  chloride  at  the  temperature  of  the  electric  arc : 

2  BC1  +  3  H  =  2  B  +  6  HC1 


352  GENERAL  CHEMISTRY 

So  prepared  the  element  is  gray  and  shows  no  crystalline  structure. 
It  melts  at  something  above  2000°,  but  has  a  considerable  vapor 
pressure  as  low  as  1600°.  In  hardness  it  closely  approaches  the 
diamond.  It  is  remarkable  for  the  extraordinary  extent  to  which  its 
electrical  resistance  falls  off  with  rise  in  temperature. 

Compounds  of  boron.  While  many  compounds  of  boron  have  been 
made,  the  oxide  and  its  derivatives  are  the  most  important.  The  only 
oxide  definitely  known  is  boric  oxide  (B2Og).  This  is  obtained  by 
heating  the  hydroxide  to  fusion,  in  the  form  of  a  transparent  glass. 
By  the  partial  reduction  of  boric  oxide  a  product  is  obtained  called 
boron  suboxide.  It  is  used  in  casting  copper,  as  its  addition  to  the 
melted  metal  prevents  the  formation  of  blowholes  on  cooling. 

The  acids  of  boron.  The  hydroxide  B(OH)3  is  known  as  boric,  or 
boracic,  acid,  but  it  has  almost  no  true  acid  properties  and  forms  no 
salts.  It  is  made  by  treating  a  hot  solution  of  borax  with  sulfuric 
acid  and  cooling,  when  the  boric  acid  crystallizes  in  shining  crystals, 
which  are  greasy  to  the  touch,  like  talc: 

Na2B40,  +  5  H20  +  H2SO4  =  Na2SO4  +  4  B(OH)3 

Its  most  remarkable  physical  property  is  that  it  is  volatile  with  steam, 
which  is  very  unusual  in  such  a  compound.  It  has  mild  antiseptic 
properties  and  is  employed  in  surgery  and  to  a  slight  extent  as  a 
preservative. 

When  boric  acid  is  carefully  heated,  it  first  forms  metaboric  acid 
(HBO2),  and  this  in  turn,  at  a  higher  temperature,  forms  tetraboric 

B(OH)3 


Salts  of  boric  acids.  While  salts  of  metaboric  acid  are  known,  most 
of  the  borates  are  derived  from  condensed  acids,  as  is  indicated  by  the 
formulas  for  borax  (Na2B4O7)  and  colemanite  (Ca2B6On). 

Borax,  which  sometimes  has  5  molecules  of  water  of  crystallization 
and  sometimes  10,  is  found  abundantly  in  nature  in  certain  localities. 
In  the  United  States  the  borax  of  commerce  is  all  manufactured  by 
digesting  colemanite  with  sodium  carbonate  : 

2  Ca2B6°n  +  4  Na2C03  +  H20  =  4  CaC03  +  3  Na2B4O7  +  2  NaOH 
It  is  also  formed  when  boric  acid  is  treated  with  sodium  hydroxide  : 
4  B(OH)3  +  2  NaOH  =  Na2B4O7  +  7  H2O 


THE   SILICON  FAMILY  AND  BOKON  353 

It  is  a  colorless  well-crystallized  salt.  In  solution  it  is  strongly 
hydrolyzed  and  has  an  alkaline  reaction.  It  is  much  used  in  place  of 
soap  in  household  cleaning  and  for  softening  hard  water. 

When  the  crystallized  salt  is  heated,  it  swells  up  to  a  sort  of  froth, 
loses  its  water  of  crystallization,  and  then  fuses  to  a  clear  glass. 
This  glass  readily  dissolves  various  metallic  oxides  and  acquires 
characteristic  colors  from  them.  This  property  is  turned  to  account 
in  chemical  analysis  in  testing  for  the  presence  of  certain  metals. 
Owing  to  this  same  property  it  is  used  as  a  flux  in  brazing  or  hard 
soldering.  It  dissolves  all  the  metallic  oxides  which  coat  the  surfaces 
of  the  metals  to  be  joined,  and  produces  a  clean  surface  to  which  the 
solder  will  adhere. 

If  it  is  remembered  that  borax  contains  an  excess  of  acid  anhydride, 
which  will  be  more  readily  seen  if  its  formula  is  written  2  NaBO2  •  B2O3, 
it  is  easy  to  understand  its  solvent  action  on  metallic  oxides.  They 
unite  with  this  excess  of  acid,  as  indicated  in  the  equation 

2  NaBO2-  B2O3  +  CuO  =  2  NaBO2-  Cu(BO2)2 

The  polysilicates  would  serve  the  same  purpose  if  they  melted  at  a 
sufficiently  low  temperature.  Some  of  the  glasses  formed  in  this  way 
are  useful  in  the  process  of  covering  ironware  with  enamels,  since 
they  are  well  adapted  to  binding  the  enamel  to  the  iron,  and  they  are 
often  an  ingredient  of  the  enamel  itself.  This  use  constitutes  one  of 
the  chief  applications  of  borax. 


CHAPTER  XXYI 

THE  PHOSPHORUS  FAMILY 


ATOMIC 
WEIGHT 

DENSITY 
OF  SOLID 

MELTING 
POINT 

BOILING 
POINT 

COLOR 

Nitrogen 

14  01 

10265 

—  2105 

—  195  7 

Snow-white 

Phosphorus  (yellow)   . 

31.04 

1.83 

44.1 

287. 

Pale  yellow 

Arsenic  (gray)  .     .     . 

74.96 

5.73 

volatilizes 

450. 

Dull  gray 

Antimony      .... 

120.2 

6.62 

630. 

1440. 

Silver-white 

Bismuth    

208.0 

9.80 

269. 

1420. 

Reddish-white 

The  group.  The  elements  of  Group  V  have  a  maximum  valence 
of  5.  The  type  element  nitrogen  is  followed  by  a  family  consisting 
of  phosphorus,  arsenic,  antimony,  and  bismuth.  The  acid-forming 
qualities,  which  are  very  strong  in  phosphorus,  gradually  decline  with 
increasing  atomic  weight,  until  in  bismuth  they  have  almost  entirely 
disappeared  and  are  replaced  by  those  of  a  pronounced  metal.  Along 
with  this  change  in  chemical  conduct  there  is  a  more  or  less  regular 
variation  in  all  physical  constants,  as  is  indicated  in  the  table. 

The  other  family  in  the  group,  comprising  vanadium,  columbium, 
and  tantalum,  is  not  so  well  known.  It  will  be  convenient  to  describe 
them  later  on  in  connection  with  certain  other  rare  elements. 

PHOSPHORUS 

Historical.  The  isolation  of  phosphorus,  which  preceded  the  recog- 
nition of  any  of  its  compounds,  was  accomplished  by  the  alchemist 
Brandt  of  Hamburg  in  1669.  He  obtained  it  by  distilling  the  resi- 
dues from  evaporated  urine.  Its  preparation  was  long  guarded  as  a 
secret,  and  the  substance  commanded  a  great  price  until  about  1770, 

when  Scheele  and  Gahn  demonstrated  its  occurrence  in  bone  ash  and 

• 

prepared  it  from  this  source. 

Occurrence.  Phosphorus  occurs  in  nature  exclusively  as  derivatives 
of  phosphoric  acid  (H3PO4).  Of  these,  apatite  (3  Ca3(PO4)2-CaF2)  is 
the  only  one  found  in  primitive  rock ;  minute  crystals  of  it  are  also 
present  in  most  soils.  The  other  phosphates  are  either  derived  from 

354 


THE  PHOSPHORUS   FAMILY 


355 


FIG.  128 


fossil  remains,  as  is  true  of  the  great  beds  of  phosphate  rock  known 
as  phosphorite,  which  is  essentially  calcium  phosphate  (Ca3(PO4)2), 
or  are  of  sedimentary  origin,  as  in  the 
case  of  the  iron  phosphate  associated 
with  bog  iron  ore. 

The  mineral  portion  of  bones  is 
largely  calcium  phosphate,  and  organic 
phosphates  appear  to  be  essential  con- 
stituents of  nerve  tissue  as  well  as  of 
the  germs  of  seeds. 

Preparation.  At  present  phosphorus 
is  usually  prepared  by  heating  min- 
eral phosphates  or  bone  ash  with  coke 
and  sand,  in  a  specially  constructed 
electrical  furnace,  the  main  features 
of  which  are  shown  in  Fig.  128.  At 

the  temperature   employed  silica   is  not  volatile,   while  phosphoric 
anhydride  (P2O5)  is.    The  volatile  anhydride  is  displaced: 

Ca8(P°4)2  +  3  S1°2  -  3  CaSi°3  '+  P2°5 

The  carbon  present  then  reduces  the  oxide  of  phosphorus : 
2P2O5+10C=10CO+P4 

The  charge  of  phosphate,  coke,  and  sand  is  fed  into  the  furnace  from  the  hop- 
per A  by  the  worm  screw  B,  while  the  liquid  calcium  silicate  is  drawn  off  at  C 
as  a  slag  or  glass.  The  vapors  of  phosphorus  escape  at  D,  together  with  the  car- 
bon monoxide,  and  are  condensed  to  a  liquid  by  a  suitable  water  condenser.  The 
liquid  is  purified  by  filtration  through  chamois  skin  or  porous  stone,  and  is  finally 
run  into  molds  immersed  in  cold  water,  and  hardened  into  sticks  for  the  market. 

White,  or  yellow,  phosphorus.  The  phosphorus  prepared  in  this  way 
is  called  white,  or  yellow,  phosphorus.  It  is  really  a  colorless,  trans- 
lucent, waxy  solid  which  melts  at  44.1°,  boils  at  287°,  and  has  a 
density  of  1.83  at  ordinary  temperatures.  It  is  insoluble  in  water 
but  dissolves  readily  in  many  solvents,  such  as  carbon  disulfide  and 
turpentine,  crystallizing  from  them  in  the  regular  system.  It  is  an 
extremely  active  substance,  combining  directly  with  most  of  the  ele- 
ments, especially  with  oxygen.  When  very  slightly  warmed  in  the 
air  it  takes  fire  and  burns  with  a  sputtering  flame,  which  becomes 
very  brilliant  in  an  atmosphere  of  oxygen.  The  product  of  combus- 
tion is  the  oxide  P2O5-  The  warmth  of  the  fingers  is  sufficient  to 


356  GENERAL  CHEMISTRY 

bring  the  substance  to  its  kindling  temperature,  so  that  the  greatest 
care  is  necessary  in  handling  it.  It  is  always  preserved  and  handled 
under  water.  It  is  a  violent  poison. 

The  slow  combustion  of  phosphorus.  There  are  many  peculiarities  attending  the 
slow  combustion  of  phosphorus.  A  stick  lying  half  covered  with  water  is  slowly 
oxidized,  and  ozone,  recognizable  by  its  odor  and  its  action  on  starch  iodide  paper, 
is  formed  at  the  same  time.  Of  the  total  oxygen  used  up,  half  goes  to  oxidize 
the  phosphorus  and  half  to  combine  with  molecular  oxygen  to  form  ozone.  The 
energy  absorbed  in  the  formation  of  the  latter  is  furnished  by  the  combustion  of 
the  phosphorus. 

In  the  slow  oxidation  of  phosphorus,  light  (which  can  be  seen  in  a  dark  room) 
is  given  off,  though  the  temperature  is  far  too  low  for  incandescence.  This 
striking  phenomenon  suggested  the  name  phosphorus,  "  light  bearer,"  and  although 
other  bodies  act  in  the  same  way,  the  phenomenon  is  called  phosphorescence. 
In  an  atmosphere  of  pure  oxygen  there  is  no  phosphorescence.  It  is  only  when 
the  partial  pressure  of  the  oxygen  falls  below  320  mm.  at  0°  that  it  appears. 

Red  phosphorus.  When  white  phosphorus  is  heated  out  of  contact 
with  oxygen,  it  is  converted  into  quite  a  different  body,  called  red 
phosphorus.  This  change  goes  on  very  slowly  at  ordinary  tempera- 
tures, is  very  marked  at  about  250°,  and  at  300°  is  very  rapid.  It  can 
be  hastened  by  catalyzers,  such  as  iodine  or  selenium.  If  the  red  phos- 
phorus is  vaporized,  it  gives  a  vapor  which  is  identical  with  that  from 
the  white  form,  and  on  condensation  the  latter  variety  is  always  ob- 
tained. Measurements  of  the  vapor  density  show  that  the  vapor  has 
the  formula  P4  (p.  318). 

Red  phosphorus  is  microcrystalline,  is  not  soluble  in  any  solvent,  is 
not  poisonous,  and  has  a  smaller  heat  of  combustion  than  the  white 
form.  It  does  not  combine  with  oxygen  at  ordinary  temperatures  and 
can  be  handled  with  entire  safety.  It  has  no  constant  physical  proper- 
ties, varying  from  scarlet  to  purple-red  in  color,  and  from  2.10  to  2.30 
in  density.  Its  heat  of  combustion  is  also  variable.  It  is  therefore 
not  a  homogeneous  substance. 

A  different  form,  called  metallic  phosphorus,  obtained  by  crystal- 
lizing phosphorus  from  lead,  has  long  been  known.  It  resembles  iodine 
in  appearance.  The  researches  of  Cohen  indicate  that  red  phosphorus 
is  really  a  solution  of  the  white  form  in  the  metallic,  these  latter  two 
being  the  only  true  modifications. 

Matches.  Aside  from  small  uses  in  the  laboratory,  phosphorus  is  employed 
almost  exclusively  in  the  manufacture  of  matches.  Friction  matches  containing 
phosphorus  first  came  into  use  in  1827,  and  at  present  two  general  varieties 
are  in  common  use.  The  more  common  variety  is  made  by  dipping  the  match 


THE  PHOSPHORUS  FAMILY  357 

stick  first  into  some  inflammable  substance,  such  as  melted  paraffin,  and  after- 
wards into  a  paste  consisting  of  (1)  white  phosphorus  or  phosphorus  sesqui- 
sulfide,  P4S3,  (2)  some  oxidizing  substance,  such  as  manganese  dioxide,  red  lead, 
or  potassium  chlorate,  and  (3)  a  binding  material,  such  as  glue  or  dextrin.  On 
friction  the  phosphorus  is  ignited,  the  combustion  being  supported  by  the  oxidiz- 
ing agent  and  communicated  to  the  wood  by  the  burning  paraffin.  In  sulfur 
matches  the  paraffin  is  replaced  by  sulfur. 

In  the  Swedish,  or  safety,  match,  red  phosphorus,  an  oxidizing  agent,  and  some 
gritty  material,  such  as  powdered  glass,  are  mixed  with  glue  and  placed  on  the 
side  of  the  box.  The  match  tip  is  provided  with  an  oxidizing  agent  and  an  easily 
combustible  substance,  usually  antimony  sulfide.  The  match  cannot  be  ignited 
easily  by  friction  except  on  the  prepared  surface. 

Constant  working  with  white  phosphorus -frequently  results  in  dreadful  dis- 
eases of  the  bones  of  the  face,  while  many  disastrous  fires  are  caused  by  the 
accidental  ignition  of  the  ordinary  match.  On  both  accounts  the  manufacture 
and  use  of  such  matches  is  prohibited  by  law  in  many  countries.  The  Congress 
of  the  United  States  has  accomplished  the  same  end  by  imposing  a  prohibitive 
tax  upon  white  phosphorus  matches  (two  cents  per  hundred  matches),  the  tax 
to  take  effect,  July,  1913.  Both  the  export  and  import  of  such  matches  is  also 
prohibited.  After  that  date  all  manufacturers  will  substitute  the  sulfide  P4S3 
for  white  phosphorus. 

Phosphides.  Phosphorus  combines  directly  with  the  great  majority 
of  the  elements  forming  phosphides.  Many  of  the  metallic  phosphides 
can  be  obtained  by  reducing  the  corresponding  phosphate  with  carbon : 

Ca3(P°4)2  +  8  C  =  Ca3P2  +  8  C° 

Some  iron  ores,  especially  limonite,  contain  iron  phosphate,  and  the  pro- 
duction of  iron  phosphide  during  the  reduction  of  the  ore  is  a  source 
of  much  trouble  in  steel  making,  since  it  is  apt  to  remain  dissolved  in 
the  metal. 

Phosphides  of  hydrogen ;  the  phosphines.  Phosphorus  forms  three 
compounds  of  hydrogen,  known  as  the  phosphines.  They  are  PH8, 
a  gas ;  P2H4,  a  colorless  liquid ;  and  P12H6,  a  yellow,  flocculent  solid. 
Of  these,  gaseous  phosphine  is  the  best  known  and  is  the  one  referred 
to  when  the  term  phosphine  is  used.  Like  the  hydrogen  compounds 
of  most  of  the  nonmetals,  it  can  be  prepared  by  decomposing  a  metallic 
phosphide  by  water  or  an  acid : 

Ca3P2  +  6  H20  =  3  Ca(OH)2  +  2  PH3 

It  is  more  easily  prepared  from  familiar  reagents  by  boiling  white 
phosphorus  with  a  concentrated  solution  of  a  strong  base.  The  reac- 
tion is  rather  complicated,  the  equation  being 

P4  +  3  KOH  +  3  H2O  =  PH3  +  3  KH2PO2 


358 


GENERAL  CHEMISTRY 


The  compound  may  be  conveniently  prepared  by  the  use  of  the  apparatus 
illustrated  in  Fig.  129.  A  200-cc.  flask  A  is  half  filled  with  a  concentrated  solu- 
tion of  potassium  hydroxide,  a  few  small  pieces  of  phosphorus  dropped  in,  and 
the  air  in  the  flask  displaced  by  a  current  of  coal  gas  admitted  through  B.  When 
the  contents  of  the  flask  is  heated,  phosphine  is  formed,  and,  passing  through  the 

delivery   tube,   is    liberated   just 

B  under  the  surface  of    the  water 

in  C.  As  each  bubble  makes  its 
escape  into  the  air  it  takes  fire 
spontaneously  and  forms  a  ring  of 
white  smoke,  consisting  of  phos- 
phorus pentoxide  (P2O5). 

The  gas  has  a  disagree- 
able odor  and  is  very  poison- 
ous. It  boils  at  -  86.2°.  It 
burns  readily  but  is  not  spon- 
taneously inflammable.  In  its 
preparation  some  of  the  liquid 
phosphine  is  formed,  which, 
being  spontaneously  inflam- 
mable, ignites  the  gas.  The  latter  may  be  freed  from  the  liquid  by 
bubbling  it  through  alcohol,  in  which  the  liquid  phosphine  is  soluble, 
or  by  conducting  it  through  a  freezing  mixture  which  condenses  the 
higher  boiling  product  to  the  liquid  state. 

In  formula,  gaseous  phosphine  resembles  ammonia,  but  in  the  pres- 
ence of  water  it  has  no  basic  properties.  When  dry  it  will  combine 
with  the  hydrides  of  the  halogens,  forming  compounds  which  in  for- 
mula resemble  ammonium  salts,  and  which  on  this  account  are  called 
phosphonium  salts.  The  best  known  of  these  is  the  iodide,  which  sub- 
limes in  colorless,  glittering  crystals : 


FIG. 129 


PH  +HI 


PH4I 


The  action  is  reversible,  and  in  the  air  the  solid  rapidly  dissociates  into 
its  components.  The  chloride  and  bromide  are  formed  only  under  in- 
creased pressure.  It  will  be  shown  in  a  later  chapter  (p.  411)  that  the 
ammonium  halides  dissociate  in  the  same  way  when  heated,  but  that  the 
temperature  at  which  dissociation  becomes  perceptible  is  much  higher. 
Liquid  phosphine  (P2H4)  is  colorless  and  strongly  refracts  light.  It 
boils  at  57°  and  has  about  the  same  density  as  water.  It  is  analogous 
in  formula  to  hydrazine  (N2H4),  and  the  two  compounds  have  many 
analogous  derivatives,  although  phosphine  has  no  basic  prdfferties. 


THE  PHOSPHOEUS  FAMILY 


359 


Halogen  compounds.  Phosphorus  combines  with  the  several  halo- 
gens directly  and  with  the  greatest  ease,  forming  two  series  of  com- 
pounds. In  the  one  it  is  trivalent  and  in  the  other,  with  the  exception 
of  the  iodide,  pentavalent.  A  table  of  these  compounds  will  indicate 
their  characteristics : 

THE  HALOGEN  COMPOUNDS  OF  PHOSPHORUS 


MELTING 
POINT 

BOILING 
POINT 

HEAT  OF 
FORMATION 

Trifluoride  .     . 

.     PFS 

-160. 

-95 

108,000  cal. 

Pentafluoride    . 

.     PF5 

-83. 

-75 

Trichloride  .     . 

.     PC13 

<-115. 

76 

76,000  cal. 

Pentachloride  . 

•       PC15 

148. 

140 

107,000  cal. 

Tribromide  .     . 

.     PBr3 

41.5 

175 

44,800  cal. 

Pentabromide  . 

•     PBr5 

decomp. 

63,000  cal. 

Tetra-iodide      . 

•        *V4 

110. 

decomp. 

19,800  cal. 

Tri-iodide     .     . 

•     "s 

60. 

decomp. 

10,900  cal. 

It  will  be  sufficient  to  describe  the  compounds  with  chlorine,  as 
they  are  the  best  known  and  have  important  uses.  The  others  are 
made  by  analogous  methods  and  have  similar  chemical  characteristics. 

Phosphorus  trichloride  (PC13).  This  liquid  is  obtained  by  passing 
a  current  of  dry  chlorine  over  white  phosphorus,  the  two  elements 
combining  with  a  flame: 

P4  +  6C12  =  4PC13 

It  is  a  colorless  liquid,  of  density  1.6,  which  fumes  strongly  in  the  air, 
owing  to  the  action  of  moisture  upon  it : 

PC13  +  3  HOH  =  P(OH)3  +  3  HC1 

A  similar  reaction  takes  place  with  the  majority  of  compounds  con- 
taining a  hydroxyl  group.  For  this  reason  the  trichloride  is  a  valuable 
reagent  for  determining  the  presence  of  hydroxyl  groups  in  compounds, 
especially  those  of  carbon.  The  reaction  with  alcohol  will  illustrate 
this  property : 

PC13  +  3  C2H5OH  =  P(OH)3  +  3  C2H5C1 

Phosphorus  pentachloride  (PC15).  The  solid,  pale  yellow  pentachlo- 
ride  is  obtained  by  passing  chlorine  into  the  trichloride,  the  reaction 
being  attended  by  the  evolution  of  considerable  heat: 


360  GENERAL  CHEMISTEY 

When  the  solid  pentachloride  is  heated  to  a  moderately  high  tempera- 
ture, this  action  is  reversed  and  dissociation  occurs,  just  as  in  the  case  of 
ammonium  chloride.  The  melting  point  can  be  determined  only  when 
the  substance  is  confined  in  a  sealed  tube,  which  prevents  the  vapors 
from  escaping.  Like  the  trichloride,  this  substance  acts  upon  water 
and  other  hydroxyl  compounds.  With  excess  of  water  the  reaction  is 

PC15  +  4  H20  =  PO(OH)S  +  5  HC1 

When  insufficient  water  is  present,  a  liquid  oxychloride  is  formed, 
named  phosphoryl  chloride: 

PC15  +  H2O  =  POC18  +  2  HC1 

Oxides  of  phosphorus.  Three  oxides  of  phosphorus  are  known, 
namely,  P2O5,  P2O4,  and  P4Og.  Vapor-density  determinations  show 
that  the  latter  has  the  double  formula  P.O.,  but  it  is  more  convenient 

4       ti 

to  use  the  simpler  formula  P2O3,  especially  since  the  molecular  weight 
of  the  other  oxides  is  not  known,  and  the  compound  is  usually  called 
the  trioxide.  Moreover,  we  have  no  knowledge  as  to  the  molecular 
weight  of  any  of  these  compounds  in  the  solid  state. 

Phosphorus  pentoxide  (P206).  The  pentoxide  is  obtained  as  a  snowlike 
solid  by  the  complete  combustion  of  phosphorus.  It  volatilizes  only  at 
a  white  heat.  Its  heat  of  formation  is  very  great,  namely,  369,900  cal., 
and  accordingly  it  is  very  stable  and  has  almost  no  oxidizing  proper- 
ties. Its  most  remarkable  characteristic  is  its  activity  toward  water, 
with  which  it  combines  with  great  energy.  On  this  account  it  is  the 
most  effective  drying  agent  known,  very  considerably  surpassing 
calcium  chloride  and  sulfuric  acid.  It  will  also  abstract  the  elements 
of  water  from  many  hydroxyl  compounds.  For  example,  it  converts 
many  acids  into  anhydrides,  as  shown  in  the  equation 

2  HN03  +  P205  =  N205  +  2  HP08 

Phosphorus  trioxide  (P203  or  P406).  This  oxide  is  formed  by  burn- 
ing phosphorus  under  conditions  which  do  not  provide  enough  oxy- 
gen for  complete  combustion.  It  often  occurs  as  an  impurity  in  the 
pentoxide.  This  also  is  a  snowlike  solid,  which  melts  at  22.5°  and 
boils  at  173.1°,  so  it  can  be  separated  from  the  pentoxide  by  distilla- 
tion. The  tetroxide  P2O4,  a  colorless  crystalline  solid,  is  formed  by 
heating  the  trioxide : 


THE  PHOSPHORUS  FAMILY 


361 


The  acids  of  phosphorus.  The  pentoxide  and  trioxide  are  both  typi- 
cal acid  anhydrides.  They  combine  with  water  directly  and,  like  the 
oxide  of  silicon,  each  gives  rise  to  a  number  of  acids  which  differ  from 
each  other  not  in  the  valence  of  the  phosphorus  but  in  the  ratio  be- 
tween the  oxide  and  the  water.  Their  formulas  and  relations  are 
shown  in  the  following  table: 


PHOSPHORIC  ACIDS 

PHOSPHOROUS  ACIDS 

Meta-    . 

•   P2o6- 

h  H20  =  2 

HP03 

Meta-    . 

•     PAH 

h  H20  =  2 

HPO2 

Pyro-     . 

•        P2°5  - 

h  2  H20  = 

H4P207 

Pyro-     . 

.     P203i 

h  2  H20  - 

H4P205 

Ortho-  . 

•    PA- 

h  3  H20  = 

2  H3PO4 

Ortho-  . 

•    P,(V 

h  3  H20  = 

2  H3PO3 

The  phosphoric  acids.  Since  all  three  of  the  phosphoric  acids  are 
derived  from  the  stable  pentoxide,  it  is  not  surprising  that  they  share 
its  stability  as  regards  reduction. 

1.  Orthophosphoric  acid  (H3POJ.  This  acid,  representing  the  greatest 
degree  of  hydration  of  the  stable  oxide,  is  the  form  into  which  all 
other  acids  of  phosphorus  tend  to  pass  when  in  solution.  Pure  hydro- 
gen phosphate  is  prepared  by  oxidizing  white  phosphorus  with  nitric 
acid,  evaporating  the  solution  to  a  sirup,  and  cooling.  The  compound 
separates  in  colorless  crystals,  which  melt  at  40°.  A  commercial  grade 
of  acid  is  prepared  by  treating  calcium  phosphate  with  concentrated 
sulfuric  acid  and  filtering  off  the  insoluble  calcium  sulfate: 

Ca,(P04)a  +  3  H2S04  -  3  CaSO4  +  2  H3PO4 

In  solution  hydrogen  phosphate  is  an  acid  of  medium  strength  and 
is  tribasic.  It  is  therefore  capable  of  forming  three  series  of  salts, 
according  as  one,  two,  or  three  hydrogen  atoms  are  replaced  by  metals. 
The  composition  and  method  of  naming  such  salts  are  illustrated  in 
the  following  table,  which  gives  the  formulas  and  the  different  names 
of  the  three  sodium  salts: 


FORMULA 

NAME 

Xa,P04     .     .     .   ' 
Na,HP04.     .     . 
XaH2PO4  .     .     . 

Trisodium  phosphate  ;  normal,  or  tertiary,  sodium  phosphate 
Disodium  phosphate  ;  secondary  sodium  phosphate 
Monosodium  phosphate  ;  primary  sodium  phosphate 

The  heavy  metals,  such  as  silver  or  mercury,  form  only  the  normal, 
or  tertiary,  phosphates.  Mixed  salts  are  also  known,  the  most  familiar 
one  being  microcosmic  salt  (NH4NaHPO4  •  4  H2O). 


362  GENERAL  CHEMISTEY 

The  tertiary  salts  of  the  alkali  metals,  sodium  and  potassium,  are 
very  strongly  hydrolyzed  in  solution  : 

Na8PO4  +  H2O  +=±Na2HPO4  +  NaOH 

Even  the  secondary  salt  Na2HPO4,  which  is  the  common  sodium 
phosphate,  has  a  decidedly  basic  reaction  in  solution. 

2.  Pyrophosphoric  acid  (J74P207).  This  acid  can  be  obtained  by  heat- 
ing the  orthophosphoric  acid  to  213°.  It  forms  a  colorless  glassy  mass 
which  melts  at  about  61°.  Its  sodium  salt  is  prepared  by  heating 
ordinary  sodium  phosphate,  and  from  this  salt  others  are  readily 
obtained  : 


43.  Metaphosphoric  acid  (HP03).    Metaphosphoric  acid  is  obtained  by 
strongly  heating  either  the  ortho  or  the  pyro  acid  : 


It  is  the  first  product  formed  when  the  pentoxide  acts  upon  water  : 
p2°5  +  H20  =  2  HP03 

It  melts  at  a  high  temperature  and,  on  cooling,  forms  a  glass  called 
glacial  phosphoric  acid.  At  very  high  temperatures  it  is  volatile.  Its 
sodium  and  potassium  salts  can  be  obtained  by  heating  the  corre- 
sponding primary  orthophosphate  : 


While  its  simplest  formula  is  HPO3,  both  the  acid  and  its  salts  exist 
in  a  variety  of  polymeric  forms,  the  acids  having  the  general  formula 
(HPOg)w,  in  which  n  is  an  integer. 

Use  of  metaphosphates  as  fluxes.    When  metaphosphoric  acid  is  dissolved  in 
water,  it  slowly  combines  with  it,  forming  the  ortho  acid  : 


In  a  similar  way,  when  sodium  metaphosphate  is  heated  with  metallic  oxides,  it 
forms  mixed  salts  of  the  ortho  acid  : 

NaPO3  +  CuO  =  NaCuPO4 

The  salts  so  formed  remain  dissolved  in  the  excess  of  metaphosphate  and,  on 
cooling,  harden  to  a  glass.  This  frequently  acquires  a  color  which  is  character- 
istic of  the  oxide  so  dissolved,  and  suggests  a  method  of  testing  for  the  presence  of 
certain  metals  in  materials  of  unknown  composition.  A  bead  of  the  metaphos- 
phate is  first  formed  by  fusing  a  crystal  of  microcosmic  salt  on  a  loop  of  wire  in 
a  Bunsen  flame  : 

NH4NaHP04  -  NaPO3  +  NH3  +  II2O 


THE  PHOSPHORUS  FAMILY  363 

This  bead  is  then  dipped  into  the  powdered  material  to  be  examined,  and  re- 
heated. The  color,  both  while  hot  and  when  cold,  shows  the  presence  of  certain 
metals,  such  as  iron,  copper,  cobalt,  and  chromium.  It  will  be  seen  that  the 
principle  is  the  same  as  in  the  case  of  the  borax  bead  (p.  353). 

The  phosphorous  acids.  The  only  one  of  the  phosphorous  acids  at  all 
well  known  is  the  ortho  acid  HgPO3,  commonly  called  merely  phos- 
phorous acid.  It  is  best  prepared  by  treating  phosphorus  trichloride 
with  water:  pcl^  +  3  J^Q  =  P(QH)3  +  3  HC1 

It  can  be  obtained  from  solution  in  large,  transparent  crystals  melting 
at  71°.  Although  it  contains  three  atoms  of  hydrogen  in  the  molecule, 
it  is  a  dibasic  acid,  only  two  of  the  hydrogen  atoms  being  replaceable 
by  metals.  On  this  account  its  formula  is  sometimes  written  H2HPO3. 
Like  the  trioxide  from  which  it  is  derived,  it  takes  up  oxygen  very 
readily  and  is  a  strong  reducing  agent.  When  heated  by  itself  it 
undergoes  an  interesting  reaction,  in  which  one  portion  is  oxidized  to 
phosphoric  acid,  while  another  is  reduced  to  phosphine : 

4  H  PO,  =  3  H  PC)  +  PH 

oo  O  4  o 

Hypophosphorous  acid  (H3P02).  There  are  a  number  of  acids  of  phos- 
phorus which  are  not  derived  from  any  known  oxide,  the  most  im- 
portant one  being  hypophosphorous  acid.  When  white  phosphorus 
is  boiled  with  concentrated  potassium  hydroxide  (see  preparation  of 
phosphine),  the  salt  KH2PO2  is  obtained,  and  from  this  the  free  acid, 
as  well  as  other  salts,  can  be  prepared.  The  acid  crystallizes  in  beauti- 
ful transparent  crystals,  which  melt  at  17.4°.  Although  it  contains 
three  atoms  of  hydrogen  to  the  molecule,  it  is  a  monobasic  acid,  the 
formula  being  sometimes  written  H  •  H2PO2.  Both  the  acid  and  all 
its  salts  are  soluble  in  water  and  are  strong  reducing  agents.  The 
hypophosphites  find  frequent  applications  in  medicine. 

Fertilizers.  Phosphorus  appears  to  be  essential  to  both  animal  and 
vegetable  organisms  and  must,  as  a  consequence,  be  present  in  every 
fertile  soil.  Since  crops  are  constantly  removed  from  cultivated  land, 
and  since  phosphorus  compounds  are  never  very  abundant  in  it,  the 
soil  is  gradually  exhausted  of  this  element,  the  supply  of  which  must 
be  renewed  in  some  way.  Animal  manures  are  of  the  greatest  value 
not  only  for  phosphorus  but  also  for  nitrogen  compounds,  but  the 
supply  is  never  sufficient.  The  large  deposits  of  rock  phosphate  found 
in  various  parts  of  the  country,  as  in  Florida,  Tennessee,  and  Dakota, 
are  the  only  source  from  which  an  adequate  supply  can  be  derived. 


364  GENERAL  CHEMISTRY 

THe  pulverized  rock  is  sometimes  applied  directly,  but  usually  it  is 
first  treated  with  sulfuric  acid,  when  one  or  both  of  the  reactions 
expressed  in  the  following  equations  takes  place : 

Ca3(P04)2  +  H2S04  =  2  CaHP04  +  CaSO4 
Ca3(P04)2  +  2  H2S04  =  Ca(H2P04)2  +  2  CaSO4 

This  treatment  appears  to  greatly  increase  the  value  of  the  rock  as  a 
fertilizer.  Doubtless  this  is  partly  due  to  the  increase  in  the  solubility 
of  the  phosphorus  compounds,  for  the  normal  calcium  phosphate  is 
practically  insoluble  in  water,  while  the  primary  salt  Ca(H2PO4)2  is 
freely  soluble,  and  the  secondary  salt  CaHPO4,  though  insoluble  in 
water,  is  soluble  in  weak  organic  acids,  such  as  are  present  about  the 
roots  of  growing  plants. 

When,  however,  these  compounds  are  spread  upon  the  soil,  which  usually  con- 
tains oxides  of  basic  character,  as  well  as  carbonates,  the  normal  salt  must  be 
formed  once  more.  This  reaction,  whether  taking  place  in  the  soil  or  in  the 
stored  fertilizer,  is  called  reversion.  The  reaction  with  limestone  will  illustrate 
this  reversion: 

2  CaHPO4  +  CaCO3  =  Ca3(PO4)2  +  H2O  +  CO2 

In  reversion  the  normal  salt  is  precipitated  in  a  very  fine  state  of  division,  and 
is  much  better  distributed  throughout  the  soil  than  could  be  accomplished  by 
mechanical  means,  which  may  partly  explain  the  value  of  acid  treatment.  It  is 
also  true  that  calcium  sulfate  is  beneficial  to  some  soils,  since  sulfur  must  be 
present  in  small  amounts.  Care  must  be  exercised  to  avoid  excess  of  sulfuric 
acid  in  the  treatment  of  the  rock,  since  this  injures  many  soils.  It  is  clear  that 
a  given  soil  may  cease  to  be  productive  through  many  different  causes,  and  there 
are  many  cases  in  which  the  application  of  a  phosphate  fertilizer  would  be  of 
no  value  whatever. 

ARSENIC 

History  and  occurrence.  Compounds  of  arsenic  have  been  known 
from  the  earliest  times.  The  highly  colored  sulfides,  realgar  (As2S2) 
and  orpiment  (As2S8),  are  found  in  nature,  and  have  been  used  as 
pigments  since  the  time  of  Aristotle.  Arsenopyrite  (FeAsS),  known 
as  mispickel,  is  an  abundant  mineral,  and  the  corresponding  cobaltite 
(CoAsS)  is  not  rare.  The  element  is  occasionally  found  in  the  free 
form,  and  also  as  the  oxide  As2O8,  called  arsenolite.  It  is  also  very 
widely  distributed  in  traces  throughout  the  sulfide  ores  of  many 
metals,  and  these  metals,  as  well  as  products  prepared  from  the 
sulfur  derived  from  the  ores,  are  apt  to  be  contaminated  with  arsenic. 
This  is  particularly  true  of  sulfuric  acid  made  from  pyrites,  and  of 
all  materials  prepared  by  the  use  of  such  acid. 


THE  PHOSPHORUS  FAMILY  365 

Preparation.  The  element  is  prepared  by  subliming  the  natural 
product  or  by  heating  mispickel  : 

4  FeAsS  =  4  FeS  +  As4 

It  is  prepared  in  pure  form  by  reducing  the  oxide  with  carbon,  the 
arsenic  being  easily  volatile  and  condensing  again  on  a  cold  surface. 
Like  phosphorus,  it  occurs  in  several  distinct  forms. 

1.  Yellow  arsenic.   This  form  is  obtained  by  very  suddenly  cooling 
the  vapors   of  arsenic.    It  is  a  yellow  crystalline  mass  resembling 
flowers  of  sulfur,  is  soluble  in  carbon  disulfide,  and,  like  white  phos- 
phorus, crystallizes  in  the  regular  system.    Its  odor  recalls  that  of 
garlic.    It  passes  with  the  greatest  ease  into  the  gray  form. 

2.  Black  arsenic.  When  arsenic  vapor  is  cooled  more  slowly,  the  ele- 
ment condenses  on  a  cold  surface  as  a  black  mirror,  brown  in  thin  films. 

3.  Gray,  or  metallic,  arsenic.  When  the  element  is  prepared  in  quan- 
tity and  is  rather  slowly  cooled,  it  deposits  as  a  gray  crystalline  mass 
somewhat  resembling  coke.    This  is  the  form  into  which  the  others 
tend  to  pass.    It  sublimes  very  easily,  reaching  a  vapor  pressure  of 
760  mm.  at  something  over  360°  and  much  below  its  boiling  point. 
It  is  very  brittle  and  has  a  density  of  5.73.    All  forms  give  the  same 
yellowish  vapor,  the  density  of  which  corresponds  to  the  formula  As4. 

Arsenides.  Arsenic  combines  directly  with  most  of  the  elements,  forming 
arsenides.  With  the  nonmetallic  elements  these  compounds  are  for  the  most  part 
decomposed  by  water.  The  metallic  arsenides  are  frequently  formed  in  the  reduc- 
tion of  metals  whose  ores  carry  some  arsenic.  They  are  usually  stable  compounds, 
are  strongly  metallic  in  appearance,  and  are  soluble  in  many  melted  metals. 
Many  of  them  may  be  obtained  by  heating  the  metal  with  arsenic  or  by  reducing 
an  arsenate  : 


Arsenic  hydride  (arsine)  (AsH3).  Arsenic  forms  only  one  hydride, 
the  colorless  gas  arsine  (AsH3).  It  is  formed  when  hydrogen  is  lib- 
erated in  contact  with  an  arsenic  compound  : 

As203  +  12  H  =  2  AsH8  +  3  H2O 

It  can  be  prepared  free  from  hydrogen  by  the  action  of  hydrochloric 
acid  upon  sodium  arsenide  : 

Na  As  +  3  HC1  =  3  NaCl  +  AsH 

3  o 

It  is  a  gas  which  liquefies  at  —55°  and  solidifies  at  —119°.  It  has  an 
odor  like  that  of  garlic  and  is  extremely  poisonous.  It  is  very  unstable 
and  burns  in  air  with  a  bluish-white  flame,  forming  arsenic  trioxide 
(As2O8)  and  water.  It  is  decomposed  into  its  elements  by  a  very 


366 


GENERAL   CHEMISTRY 


moderate  heat.  This  is  easily  demonstrated  by  conducting  the  gas 
through  a  tube  heated  at  one  point,  the  arsenic  depositing  as  a  black 
mirror  a  little  beyond  the  hot  region.  Owing  to  the  form  in  which  it 
is  deposited,  a  mere  trace  is  easily  seen.  Because  of  this  ready  decom- 
position it  is  a  strong  reducing  agent. 

Marsh's  test  for  arsenic.  The  properties  of  arsine  make  possible  the  use  of  a 
very  delicate  method  for  detecting  its  presence,  known  as  Marsh's  test.  Hydrogen 
is  generated  in  a  flask  A  (Fig.  130)  by  the  action  of  hydrochloric  acid  upon  zinc,  the 
escaping  gases  being  dried  by  passing  through  a  calcium  chloride  tube  B,  Heat 
is  applied  at  C  for  some  time,  and  if  no  mirror  forms,  the  reagents  are  free  from 

arsenic.  The  material  to 
be  tested  is  brought  into 
solution  by  the  necessary 
preliminary  treatment 
and  is  introduced  through 
the  funnel  tube,  the  pres- 
ence of  arsenic  being  in- 
dicated by  the  speedy 
formation  of  a  mirror. 
Instead  of  heating  the 
tube  a  cold  porcelain  dish 
may  be  held  in  the  flame 
at  D,  the  arsenic  being 
deposited  upon  it  as  a 

dark  spot.  The  mirror  is  a  very  brilliant  black,  is  brownish  at  the  edges,  is  very 
easily  volatilized  when  heated,  dissolves  readily  in  a  solution  of  sodium  hypo- 
chlorite,  and  turns  yellow  on  conducting  a  current  of  hydrogen  sulfide  through  the 
tube.  All  these  characteristics  distinguish  the  arsenic  mirror  from  a  similar  one 
produced  by  the  element  antimony.  Under  favorable  conditions  the  presence  of  a 
quantity  of  arsenic  no  greater  than  0.1  mg.  can  be  detected.  Indeed,  it  is  difficult  to 
secure  zinc  and  acid  for  generating  hydrogen  which  will  not  give  a  test  for  arsenic. 

Halogen  compounds.  The  well-known  halogen  compounds  of  arsenic 
are  derived  from  the  trivalent  element,  their  chief  physical  constants 
being  given  in  the  table.  The  pentafluoride  AsF5  and  possibly  the 
pentachloride  AsCl5  are  also  known. 

HALOGEN  COMPOUNDS  OF  ARSENIC 


-pIG 


DENSITY 

MELTING 
POINT 

BOILING 
POINT 

COLOR 

Trifluoride     .     . 

AsF3 

2.7  liquid 

-8.3 

63. 

Colorless 

Trichloride    .     . 

AsCl3 

2.2  liquid 

-18. 

130.2 

Colorless 

Tribromide    .     . 

AsBr8 

3.7  solid 

81.' 

221. 

Colorless 

Tri-iodide      .     . 

AsI8 

4.4  solid 

146. 

ap.  400. 

Red 

THE  PHOSPHORUS  FAMILY 


367 


These  compounds  are  decomposed  by  water,  but  not  so  readily  as  in 
the  case  of  the  corresponding  phosphorus  compounds.  This  is  seen 
from  the  fact  that  arsenious  oxide  (As2O3)  dissolves  in  concentrated 
hydrochloric  acid,  though  it  is  but  sparingly  soluble  in  water.  When 
the  solution  is  boiled,  some  of  the  arsenic  passes  off  with  the  vapor ; 
this  does  not  happen  when  a  solution  of  the  oxide  in  water  is  boiled. 
Arsenic  chloride  is  therefore  present  in  the  solution: 

As2O3  +  6  HC1  +=±  2  AsCl3  +  3  H2O 

Oxides  of  arsenic.  There  are  two  oxides  of  arsenic:  the  trioxide 
As2O3,  or  As4O6,  and  the  pentoxide  As2O&.  These  are  both  acid 
anhydrides. 

Arsenic  trioxide  (white  arsenic)  (As203,  or  As406).  As  a  matter 
of  convenience  the  formula  of  this  oxide  is  usually  written  As2O3, 
although  the  vapor-density  measurements  show  that  it  should  be 
As4Og.  The  oxide  is  found  in  nature  in  several  forms,  and  it  is 
obtained  as  a  by-product  in  metallurgical  processes  in  which  sulfide 
ores  are  roasted.  In  this  process  the  arsenic  is  converted  into  oxide, 
which  condenses  as  a  fine  dust  on  the  cooling  of  the  gases  formed 
in  roasting.  From  the  flue  dust  the  pure  oxide  may  be  obtained  by 
sublimation,  as  it  is  readily  volatile.  When  sublimed  with  slow  cool- 
ing, the  product  is  a  trans- 
parent amorphous  glass,  which 
gradually  becomes  opaque,  like 
porcelain,  owing  to  crystalliza- 
tion. When  condensed  rapidly, 
it  forms  a  crystalline  powder  of 
very  characteristic  appearance. 

The  ready  formation  of  this  crys- 
talline deposit  is  turned  to  practical 
account  in  testing  for  the  presence 
of -the  oxide.  The  material  to  be 
tested  is  placed  in  a  small  test  tube 
of  hard  glass  and  is  heated  in  a  Bunsen  flame.  If  arsenic  oxide  is  present,  it 
will  deposit  in  characteristic  white  crystals  upon  the  colder  walls  of  the  tube,  as 
shown  in  Fig.  131. 

The  oxide  is  not  very  soluble  in  water,  1 1.  of  water  at  ordinary  tem- 
peratures dissolving  16  g.  of  the  solid.  It  has  a  weak,  sweetish  taste 
and  is  very  poisonous ;  a  dose  of  0.06  g.  has  been  known  to  prove 
fatal.  Notwithstanding  this  fact,  the  system  can  become  accustomed 


FIG.  131 


368 


GENERAL  CHEMISTRY 


to  it  by  gradually  increasing  the  dose.  The  mountaineers  of  the  Tirol 
find  that  it  increases  their  endurance  in  mountain  climbing,  and  they 
gradually  become  accustomed  to  daily  portions  four  times  as  large  as 
a  fatal  dose  for  an  ordinary  person.  The  oxide  finds  a  limited  use  in 
medicine  and  as  a  poison  for  vermin.  It  is  used  as  a  preservative  in 
taxidermy  and  as  a  mild  reducing  agent  in  some  chemical  industries. 
It  is  also  used  in  the  manufacture  of  some  insecticides  and  as  an 
ingredient  in  glass  making. 

Arsenic  pentoxide  (As20B).  This  oxide,  a  white  amorphous  solid,  can- 
not be  obtained  by  the  direct  combustion  of  arsenic  or  of  the  trioxide, 
but  is  made  by  heating  arsenic  acid : 

2  H3As04  =  As205  +  3  H2O 

It  is  not  very  stable  toward  heat,  and  at  a  high  temperature  dissoci- 
ates into  the  trioxide  and  oxygen  ;  consequently  its  molecular  weight 
is  unknown.  It  is  a  moderately  strong  oxidizing  agent. 

Acids  of  arsenic.  Arsenic  forms  two  series  of  oxygen  acids  derived 
from  the  two  oxides  quite  similar  in  formulas  to  those  of  phosphorus. 
The  names  and  formulas  of  these  are  indicated  in  the  following  table  : 


ARSENIOUS  ACIDS 

AKSENIC  ACIDS 

Orthoarsenious  acid    .     .     . 
Pyroarsenious  acid       .     .     . 
Metarsenious  acid  .... 

H3AsO3 
H4As2O5 
HAs02 

Orthoacsenic  acid  .... 
Pyroarsenic  acid    .... 
Metarsenic  acid     .... 

H3As04 
H4As2O7 
HAsO3 

The  arsenious  acids.  When  dissolved  in  water,  arsenic  trioxide  gives 
rise  to  a  number  of  different  acids,  none  of  which  have  been  obtained 
in  pure  condition  and  all  of  which  are  very  weak.  By  suitable  means 
the  salts  of  these  acids  can  be  obtained,  those  of  sodium  and  potas- 
sium being  derived  from  the  metarsenious  acid  HAsO2,  while  those 
of  most  of  the  other  metals  are  derived  from  the  ortho  acid  H8AsO8. 
Thus  the  potassium  salt  is  KAsO2,  while  the  silver  salt  is  Ag3AsOg. 
Salts  of  the  pyro  acid  H4As2O5  are  also  known.  The  copper  salts  are 
of  importance.  Scheele's  green  (CuHAsO3)  is  used  as  a  pigment. 
Paris  green,  also  called  Schweinfurt  green,  is  used  as  a  poison  for 
insects.  It  is  a  double  salt  of  the  formula  Cu3(AsO3)2  •  Cu(C2H3O2)2. 
With  freshly  precipitated  ferric  hydroxide  —  Fe(OH)3  in  colloidal  state 
—  these  acids  form  an  insoluble  compound  which  is  not  poisonous. 
This  is  the  best  antidote  for  arsenic  poisoning.  It  should  be  prepared, 


THE  PHOSPHORUS  FAMILY  369 

when  needed,  by  treating  iron  alum  with  magnesium  or  calcium  hydrox- 
ide in  the  proportion  required  by  the  equation 

2  KFe(S04)2  +  3  Mg(OH)2  =  K2SO4  +  3  MgSO4  4-  2  Fe(OH)8 

The  arsenic  acids.  When  arsenic  trioxide  is  boiled  with  nitric  acid 
and  the  solution  is  evaporated  to  a  sirup,  crystals  of  orthoarsenic  acid 
are  obtained,  which  have  the  formula  2  HgAsO4  •  H2O.  When  these 
crystals  are  heated  from  140°  to  180°,  they  lose  water  and  form  the 
pyro  acid  H4As2O7.  When  they  are  heated  to  200°,  the  meta  acid 
HAsO8  is  formed.  All  these  acids,  when  dissolved  in  water,  are 
again  converted  into  the  ortho  acid.  Salts  of  this  latter  acid  are  not 
infrequently  found  in  nature,  and  many  have  been  prepared  in  the 
laboratory.  The  reddish-brown  silver  salt  AggAsO4  and  the  mag- 
nesium ammonium  salt  MgNH4AsO4  are  of  importance  in  analytical 
chemistry.  All  these  salts  are  quite  similar  to  the  corresponding 
phosphates.  Unlike  phosphoric  acid,  arsenic  acid  is  an  oxidizing 
agent,  as  might  be  anticipated  from  the  instability  of  its  anhydride. 

Sulfides  of  arsenic.  There  are  three  well-known  sulfides  of  arsenic  : 
namely,  As2S2,  As2S3,  and  As2S5. 

Arsenic  disulfide  (realgar)  (As2S2).  Realgar  is  found  in  nature  as 
very  beautiful  red  crystals  and  is  artificially  prepared  by  heating 
arsenic  with  sulfur  in  the  proper  proportions.  It  was  formerly  much 
used  as  a  pigment.  Vapor-density  measurements  show  that  under 
900°  it  has  the  formula  As4S4. 

Arsenic  trisulfide  (orpiment)  (As2S3).  This  substance  is  found  in 
nature  in  yellow  crystals  and  is  artificially  prepared  by  heating  a 
mixture  of  arsenic  and  sulfur.  In  the  laboratory  it  results  as  an 
amorphous  yellow  precipitate  when  hydrogen  sulfide  is  conducted 
into  an  acidulated  solution  of  an  arsenious  compound: 


If  the  solution  is  neutral  or  basic,  no  precipitate  forms,  although  the 
solution  turns  yellow.  The  sulfide  remains  suspended  in  colloidal 
form  and  is  only  coagulated  when  acids  or  certain  salts  are  added. 
It  is  insoluble  in  water  and  in  acids. 

Arsenic  pentasulfide  (As2S5).  The  pentasulfide  is  slowly  precipitated 
as  a  curdy  yellow  precipitate  when  hydrogen  sulfide  is  conducted 
into  a  solution  of  arsenic  acid  containing  hydrochloric  acid: 

2  H3As04  +  5  H2S  =  As2S6  +  8  H2O 


370  GENERAL  CHEMISTRY 

Thio  salts  of  arsenic.  When  either  the  trisulfide  or  the  pentasulfide 
is  treated  with  a  solution  of  a  soluble  sulfide,  such  as  sodium  sulfide 
(Na2S),  it  is  dissolved,  and  when  the  solution  is  evaporated,  well- 
crystallized  thio  salts  are  obtained: 

As2S8  +  3  Na2S  =  2  Na3AsS8 
As2S5  +  3  Na2S  =  2  NagAsS4 

These  salts  may  be  regarded  as  derived  from  the  corresponding 
oxygen  salts  by  the  replacement  of  oxygen  with  sulfur.  Thio  salts 
corresponding  to  both  the.  pyro  and  the  meta  acids  are  also  known. 
All  these  salts  are  usually  easier  to  obtain  in  well-crystallized  form 
than  the  oxygen  salts.  On  treatment  with  acids  they  are  decomposed, 
with  the  formation  of  the  sulfides  : 

2  Na3AsS4  +  6  HC1  =  6  NaCl  +  As2S5  +  3  H2S 

These  reactions  make  it  possible  to  separate  the  sulfides  of  arsenic 
from  those  of  other  elements  which  are  insoluble  in  the  sulfides  of 
sodium  or  ammonium. 

ANTIMONY 

Historical.  Compounds  of  antimony,  as  well  as  the  element  itself, 
have  been  known  from  the  earliest  times.  The  Chaldeans  appear  to 
have  made  ornamental  vessels  of  the  metal.  The  Chinese  employed 
preparations  of  antimony  as  drugs.  The  sulfide  SbjS3,  known  as  stib- 
nite,  was  named  stibium  by  Pliny,  the  name  antimonium  first  appear- 
ing in  the  works  of  the  alchemist  Geber.  In  the  fifteenth  century, 
compounds  of  antimony  were  widely  used  as  medicines. 

Occurrence.  To  some  extent  antimony  occurs  uncombined,  particu- 
larly in  Queensland  and  New  South  Wales.  Its  chief  ore  is  stibnite, 
which  occurs  in  considerable  deposits  in  Hungary  and  Japan.  It  is 
also  found  rather  widely  distributed  in  mineral  regions,  combined  with 
sulfur  or  oxygen  or  as  double  sulfides  with  lead  or  silver. 

Preparation.  The  element  is  prepared  by  refining  the  native  product 
or,  more  usually,  from  the  sulfide  by  melting  it  with  iron  : 


Prepared  in  this  way  it  almost  always  contains  copper,  arsenic,  and 
iron.  These  are  removed  by  remelting  with  a  little  antimony  sulfide, 
whereby  the  sulfides  of  the  other  metals  are  formed  and  float  on  top 
of  the  antimony  as  a  liquid. 


THE  PHOSPHOKUS  FAMILY  3T1 

Properties.  Like  phosphorus  and  arsenic,  antimony  exists  in  sev- 
eral distinct  forms,  including  (a)  a  yellow,  nonmetallic  form,  (£>)  a 
black,  metallic  form,  and  (<?)  the  ordinary  gray  form.  The  latter 
is  a  'silvery,  shining  metal,  very  brittle,  and  forms  crystals  which,  like 
those  of  gray  arsenic,  belong  to  the  hexagonal  system.  Its  density  is 
6.62,  its  melting  point  630°,  and  its  boiling  point  1440°.  Its  vapor 
density  indicates  that  the  vapor  is  a  mixture  of  molecules  of  the 
formulas  Sb2  and  Sb4. 

Chemical  conduct.  In  its  compounds  antimony  may  be  either  diva- 
lant,  trivalent,  or  tetravalent.  In  the  compounds  in  which  it  is  tri- 
valent  it  plays  the  part  of  a  metal ;  as  a  pentavalent  element  it  is, 
like  arsenic,  a  nonmetal.  It  stands  midway  between  arsenic,  which  is 
strictly  acid-forming,  and  bismuth,  which  is  base-forming.  Antimony 
is  not  attacked  by  dilute  acids,  being  below  hydrogen  in  the  electro- 
motive series.  Concentrated  sulfuric  acid  converts  it  into  the  sulfate 
Sb2(SO4)3,  with  liberation  of  sulfur  dioxide,  and  nitric  acid  oxidizes  it 
to  the  pentavalent  condition,  forming  the  solid  HgSbO4.  At  ordinary 
temperatures  it  is  very  little  acted  upon  by  oxygen,  but  as  a  powder 
it  is  readily  combustible.  It  acts  upon  steam,  liberating  hydrogen : 

2  Sb  +  3  H20  «=±Sb208  +  3  H2 

Antimony  hydride  (stibine)  (SbH3).  Antimony  forms  only  one 
hydride  (SbHg),  corresponding  to  phosphine  and  arsine.  It  is  ob- 
tained by  methods  analogous  to  those  employed  with  arsine,  namely, 
by  the  action  of  acids  on  metallic  antimonides,  such  as  those  of  zinc 
or  magnesium,  and  by  the  action  of  nascent  hydrogen  on  antimony 
compounds.  It  is  a  gas  which  liquefies  at  — 18°  and  solidifies  at 
—  91.5°,  and  it  has  an  odor  suggesting  that  of  hydrogen  sulfide. 
In  chemical  conduct  it  closely  resembles  arsine,  but  is  more  easily 
decomposed  by  heat  (at  150°  to  200°),  and  the  mirror  formed  in  a 
cold  tube  is  readily  distinguished  from  that  of  arsenic,  as  already 
described  (p.  366). 

Halogen  compounds  of  antimony.  Antimony  forms  an  almost  com- 
plete double  series  of  halogen  compounds,  the  pentabromide  alone 
being  unknown.  They  can  be  prepared  by  direct  union  of  the  ele- 
ments, and  in  some  cases  by  the  action  of  the  halogen  acid  upon 
antimony  or  its  oxides. 

Most  of  these  compounds  are  partially  decomposed  by  water,  the 
reaction  coming  to  an  equilibrium  before  it  is  entirely  completed. 


372  GENERAL  CHEMISTRY 

With  the  chloride  the  chief  product  of  the  hydrolysis  is  the  oxychlo- 
ride  SbOCl :          SbClg  +  JJ^Q  ^—^  sbOCl  +  2  HC1 

The  following  table  gives  the  chief  characteristics  of  these  compounds : 

THE  HALOGEN  COMPOUNDS  OF  ANTIMONY 


• 

MELTING 
POINT 

BOILING  POINT 

COLOR 

Trifluoride    .     .     . 

.     SbF3 

292. 

Colorless  crystals 

Pentafluoride    .     . 

.     SbF5 

149. 

Oily  liquid 

Trichloride  . 

.     SbCl3 

73.2 

223. 

Colorless  crystals 

Pentachloride    . 

.     SbCl5 

-6. 

140.  decomp. 

Colorless  liquid 

Tribromide  . 

.     SbBr3 

90. 

275.4 

Colorless  crystals 

Tri-iodide     .     .     . 

SbI3 

167. 

401. 

Ruby-red  crystals 

Pentiodide    .     .     . 

SbT5 

78. 

decomp. 

Dark  brown  solid 

Oxides  of  antimony.  Antimony  forms  three  oxides :  Sb2O8,  a  white 
solid ;  Sb2O4,  a  white  powder ;  Sb2O5,  a  yellow  powder. 

Antimony  trioxide  (Sb203  or  Sb406).  This  oxide  is  sometimes  found 
in  nature  and  is  prepared  by  burning  metallic  antimony  in  air  and 
subliming  the  product.  It  crystallizes  in  minute  cubes,  and  its  vapor 
density  shows  it  to  have  the  double  formula  Sb4O6,  as  in  the  case  of 
the  corresponding  oxide  of  arsenic.  It  is  white  when  cold  but  yellow 
when  heated.  It  is  insoluble  in  water  and  in  most  dilute  acids,  but 
is  soluble  in  the  halogen  acids. 

The  corresponding  hydroxide  Sb(OH)3  is  a  white,  amorphous  pre- 
cipitate which  loses  water  in  several  stages,  finally  passing  into  the 
oxide.  Suspended  in  water,  it  dissolves  both  in  acids  and  in  bases, 
showing  that  it  can  act  either  as  a  weak  base  or  as  a  weak  acid. 
A  hydroxide  possessing  such  properties  is  said  to  be  amphoteric.  In 
the  presence  of  strong  bases  it  forms  salts  analogous  to  metarsen- 
ites  derived  from  the  acid  HSbO2.  An  example  is  the  sodium  salt 
NaSbO2  *  3  H2O.  The  structure  of  this  acid  can  be  inferred  from  the 
relation  which  it  sustains  to  the  hydroxide : 

.OH 
cu  '  r»TT I_CK// *^     _L  TI  n 

o D  —  \J XI  f~-  D  D  \  /"v TT  ~r  -H-2 

XOH 

When  this  partially  dehydrated  form  is  acted  upon  by  an  acid,  it 
may  form  a  salt  of  the  general  type  Sb^pj  or  sb^^o  •  In  these  salts 
the  group  SbO  plays  the  part  of  a  univalent  radical  to  which  the 


THE  PHOSPHOKUS  FAMILY  373 

name  antimonyl  has  been  given,  the  two  whose  formulas  have  just  been 
mentioned  being  antimonyl  chloride  and  nitrate  respectively.  A  num- 
ber of  salts  are  also  known  in  which  the  antimony  acts  as  a  trivalent 
metal.  Among  these  are  the  halogen  salts  and  the  sulfate  Sb2(SO4)g. 
Antimony  pentoxide  (Sb205).  The  pentoxide  is  an  amorphous  yel- 
low powder  formed  by  carefully  heating  the  nitrate.  Above  400°  it 
decomposes  into  the  tetroxide  Sb2O4  and  oxygen.  From  it  are  derived 
three  acids  similar  to  those  of  phosphorus  and  arsenic  : 

Orthoantimonic  acid  (H3SbO4)  :  an  insoluble  white  powder 
Pyroantimonic  acid  (H4Sb2O7)  :  an  insoluble  white  powder 
Metantimonic  acid  (HSbOs)  :  an  insoluble  white  powder 

Many  salts  of  these  acids  are  known,  the  sodium  salts  being  remarkable 
for  the  fact  that  they  are  but  sparingly  soluble,  which  is  very  unusual 
for  sodium  salts.  Both  the  salts  and  the  acids  have  oxidizing  properties, 
as  would  be  expected  from  the  ease  with  which  the  oxide  decomposes. 

Sulfides  of  antimony.  Three  sulfides  of  antimony  are  known,  having 
the  formulas  Sb2S3,  Sb2S4,  Sb2S5.  The  second  is  of  little  importance. 

Antimony  trisulfide  (stibnite)  (Sb2S3).  In  nature  this  substance  is 
found  in  brilliant  black  prisms.  By  conducting  hydrogen  sulfide  into 
a  solution  of  an  antimony  salt  the  same  compound  is  formed  as  an 
orange-red  precipitate.  In  nearly  neutral  solution  this  acts  like  arsenic 
trisulfide  in  forming  a  colloidal  suspension.  It  is  more  readily  soluble 
in  concentrated  acid  than  the  corresponding  arsenic  trisulfide.  Like 
the  latter,  it  dissolves  in  alkaline  sulfides  to  form  a  thio  salt  : 


Sb2S8  +  3  Na2S  = 

These  thio  salts  are  decomposed  by  acids  reprecipitating  the  sulfide. 
Antimony  trisulfide  is  easily  combustible  and  is  used  in  the  manu- 
facture of  matches.  It  is  also  used  in  the  vulcanizing  and  coloring 
of  rubber  and  as  a  pigment. 

Antimony  pentasulfide  (Sb2S5).  This  compound  is  prepared  by  con- 
ducting hydrogen  sulfide  into  a  solution  of  an  antimonic  compound  : 

2  SbCl5  +  5  H2S  =  Sb2S6  +  10  HC1 

It  is  a  dark  orange,  amorphous  substance,  insoluble  in  most  liquids. 
It  readily  decomposes  into  the  trisulfide  and  sulfur.  Like  the  penta- 
sulfide of  arsenic,  it  dissolves  in  alkaline  sulfides  to  form  thio  salts. 
The  sodium  salt  Na3SbS4  •  9  H2O,  known  as  Schlippe's  salt,  is  readily 
obtained  in  crystalline  form. 


374  GENERAL  CHEMISTRY 

Alloys.  Many  elements,  especially  metals  and  metalloids,  when 
melted  together  are  mutually  soluble,  and  on  being  cooled  the  solu- 
tion freezes  to  a  solid  called  an  alloy.  In  the  process  of  solidification 
alloys  show  many  peculiarities  depending  upon  the  character  of  the 
components  of  the  alloy.  (1)  The  several  elements  present  may  each 
separate  in  minute  crystals,  making  a  fine-grained  solid.  (2)  The 
crystals  as  they  form  may  be  solid  solutions,  each  one  being  made  up 
of  all  the  elements  present.  (3)  Definite  compounds  may  crystallize 
from  the  solution,  together  with  crystals  of  the  several  elements.  A 
great  variety  of  properties  is  therefore  to  be  found  among  the  differ- 
ent alloys.  The  alloys  of  mercury  are  called  amalgams. 

One  of  the  chief  uses  of  antimony  is  in  the  preparation  of  alloys. 
Many  alloys  of  antimony,  like  the  element  itself,  expand  on  cooling, 
and  this  property  makes  them  valuable  for  making  castings  in  which 
fine  lines  are  to  be  reproduced,  as  in  the  casting  of  type  (see  table, 

p.  375). 

BISMUTH 

Historical.  References  to  bismuth  are  found  in  the  writings  of  the 
fifteenth  century,  but  the  element  was  confused  with  other  metals, 
such  as  tin,  antimony,  lead,  and  zinc.  Little  was  known  of  it  with 
any  accuracy  until  the  researches  of  Bergman,,  about  the  middle  of  the 
eighteenth  century. 

Occurrence  and  preparation.  Bismuth  is  a  rather  rare  element,  yet  it 
is  found  widely  distributed  in  mineral  regions  both  in  this  country  and 
in  parts  of  Europe,  South  America,  and  Australia.  It  usually  occurs 
uncombined ;  but  a  great  variety  of  rare  minerals,  especially  sulfides, 
contain  it  as  a  constituent.  The  simplest  of  these  are  bismuth  glance 
(Bi2S3)  and  bismuth  ochre  (Bi2O3). 

A  considerable  quantity  of  the  metal  is  produced  from  native  bis- 
muth by  heating  the  ore  and  drawing  off  the  liquid  bismuth.  It  is  pro- 
duced from  its  compounds  by  rather  complicated  processes,  the  essential 
features  of  which  are  the  preliminary  roasting  of  the  ore  to  produce 
the  oxide,  and  the  heating  of  this  with  carbon  and  a  suitable  flux.  It 
is  also  obtained  as  a  by-product  in  the  refining  of  lead  (p.  505). 

Properties  and  conduct.  Bismuth  is  a  silvery  metal  with  a  decidedly 
ruddy  tint.  It  is  very  crystalline,  brittle,  and  has  a  high  luster.  It  is 
difficult  to  prepare  it  in  a  perfectly  pure  condition,  and  its  physical 
constants  are  therefore  not  accurately  known.  Its  density  is  about 
9.80,  its  melting  point  269°,  and  its  boiling  point  1420°.  At  the 


THE  PHOSPHORUS  FAMILY 


375 


highest  temperature  its  vapor  density  shows  the  molecule  to  be  mona- 
tomic,  which  is  apparently  true  of  all  the  metallic  elements. 

At  ordinary  temperatures  bismuth  is  not  affected  by  the  air,  but 
when  heated  it  burns  to  form  the  trioxide  Bi2Og.  Like  antimony,  it 
decomposes  steam  in  a  reversible  reaction.  It  dissolves  in  hot,  con- 
centrated sulfuric  and  nitric  acids  to  form  bismuth  salts,  but  it  is 
very  slowly  attacked  by  hydrochloric  acid  in  the  presence  of  air.  The 
halogen  elements  combine  directly  with  it,  but  not  with  great  energy. 

Alloys.  Bismuth  is  used  chiefly  in  the  making  of  alloys.  Like  anti- 
mony, the  element,  as  well  as  its  alloys,  expands  on  cooling,  and  the 
alloys  have  low  melting  points.  The  following  table  shows  the  com- 
position of  the  chief  alloys  of  antimony  and  bismuth. 

ALLOYS  OF  ANTIMONY  AND  BISMUTH 


LEAD 

TIN 

ANTIMONY 

BISMUTH 

Type  metal  (older)    . 

50. 

25. 

25 

Type  metal  (newer)  . 

60. 

10. 

30 

Pewter 

20. 

80 

Britannia  metal    .     .     . 

90. 

7 

Copper  3 

Antifriction  metal 

(Babbitt  metal)     .     . 

65. 

17. 

17 

Fusible  metals  : 

Rose's,  m.p.  93.8°  .     . 

25. 

25. 

50 

Wood's,  m.p.  60.5°      . 

25. 

12.5 

50 

Cadmium  12.5 

Newton's,  m.p.  94.5°  . 

31.75 

18.75 

50 

Alloys  of  this  kind  are  used  in  the  manufacture  of  household  utensils 
(Britannia  metal),  of  type,  of  bearings  in  machinery  (antifriction  or 
Babbitt  metal),  of  safety  plugs  for  steam  boilers,  and  for  liquid  baths 
to  secure  a  steady  temperature. 

Compounds  of  bismuth.  In  nearly  all  of  its  compounds  bismuth 
plays  the  part  of  a  trivalent  metal,  yielding  a  series  of  salts  most  of 
which  are  colorless.  It  forms  no  acids,  as  do  the  other  members  of  the 
group.  This  is  in  accord  with  the  general  rule  that  as  the  atomic 
weight  increases  in  any  periodic  family  the  metallic  properties  become 
more  pronounced.  The  element  forms  no  hydride. 

Oxides  of  bismuth.  While  bismuth  forms  a  number  of  oxides,  includ- 
ing those  having  the  formulas  BiO,  Bi2O8,  Bi2O4,  and  Bi2O5,  the  tri- 
oxide alone  is  at  all  well  known,  and  from  it  are  derived  most  of  the 
compounds  of  the  element. 


376  GENERAL  CHEMISTRY 

Bismuth  trioxide  (Bi203).  This  is  obtained  by  burning  the  metal  or 
by  heating  the  hydroxide  or  carbonate.  It  is  a  yellowish,  crystalline 
powder  melting  at  about  820°  and  volatile  at  very  high  temperatures. 
It  dissolves  in  acids  to  form  the  corresponding  salts,  such  as  the  chlo- 
ride BiCl3  •  2  H2O  and  the  nitrate  Bi(NOg)3  •  5  H2O,  both  of  which 
are  colorless  solids. 

Hydrolysis  of  bismuth  salts.  While  the  normal  salts  are  stable  in 
solutions  containing  an  excess  of  acid,  they  are  hydrolyzed  in  water 
or  in  dilute  acids.  In  this  hydrolysis  a  basic  salt  is  formed  and  not 
the  hydroxide.  With  the  chloride  the  reaction  is  expressed  in  the 
following  equation: 

BiCl3  +  2  H20  =  Bi(OH)2Cl  +  2  HC1 

Such  basic  salts  are  nearly  insoluble  and,  as  a  rule,  lose  water : 
Bi(OH)2Cl  =  BiOCl  +  H20 

This  oxychloride  may  be  regarded  as  the  chloride  of  the  base  BiO  •  OH, 
in  which  the  group  BiO  plays  the  part  of  a  univalent  radical.  This 
radical  is  called  bismuthyl,  just  as  the  corresponding  antimony  radical 
is  called  antimonyl.  There  are,  therefore,  two  series  of  salts:  the 
normal  salts,  corresponding  to  the  formulas  BiClg  and  Bi(NO3)8,  and 
the  bismuthyl  salts,  such  as  the  chloride  BiO  •  Cl,  the  nitrate  BiO  •  NO3, 
and  the  carbonate  (BiO)2COg.  A  number  of  these  bismuthyl  salts, 
particularly  the  nitrate  and  the  carbonate,  are  used  in  medicine.  They 
are  usually  called  bismuth  subnitrate  and  bismuth  subcarbonate. 

Bismuth  sulfide  (Bi2S3).  The  chief  sulfide  of  bismuth  has  the  for- 
mula Bi2Sg.  It  can  be  prepared  by  direct  union  of  the  elements  or 
by  conducting  hydrogen  sulfide  into  a  solution  of  a  bismuth  salt: 

2  BiOCl  +  3  H2S  =  Bi2S8  +  2  H2O  +  2  HC1 

Prepared  in  this  latter  way  it  is  a  black,  amorphous  precipitate, 
insoluble  in  water,  in  dilute  acids,  and  in  the  sulfides  of  sodium 
or  ammonium. 


CHAPTER  XXVII 

THE  HYDROXIDES  AND  THEIR  REACTIONS 

Before  entering  upon  a  study  of  the  metallic  elements  it  will  be  of 
advantage  to  bring  together  in  one  place  a  number  of  facts,  already 
presented  in  various  connections,  with  regard  to  the  hydroxides  of 
the  elements  and  their  chief  reactions. 

The  hydroxides  and  the  oxides.  Almost  every  element  forms  a 
hydroxide  as  one  of  its  most  important  compounds.  If  the  element  is 
capable  of  exerting  more  than  one  valence,  there  may  be  a  hydroxide 
corresponding  to  each.  By  loss  of  water  these  hydroxides  may  be 
expected  to  yield  oxides,  and  in  the  great  majority  of  cases  these 
oxides  are  well-known  bodies. 

Normal  acids.  The  hydroxide  of  an  acid-forming  element  is  some- 
times called  a  normal  acid.  In  some  cases  these  are  well-known  com- 
pounds, as  in  the  case  of  phosphorous  acid  (P(OH)3).  In  other  cases 
the  hydroxide  loses  water  so  readily  that  it  cannot  be  isolated  in  pure 
form.  For  example,  the  hydroxides  of  nitrogen  and  sulfur,  which 
should  have  the  formulas  N(OH)5  and  S(OH)6  lose  water  and  form 
the  ordinary  acids  HNO8  and  H2SQ4.  It  is  often  of  advantage,  how- 
ever, to  consider  these  latter  acids  as  dehydrated  forms  of  the  normal 
acids  N(OH)5  and  S(OH)6. 

Dehydration  involves  no  change  in  valence.  The  relation  between 
a  hydroxide  and  an  oxide  is  shown  in  the  following  equations,  E 
standing  for  any  element: 

2  E(OH)  ^z±  E2O  +  H2O 


These  equations  indicate  the  fact  that  in  the  process  of  dehydration 
there  is  no  change  in  valence  on  the  part  of  the  element,  and  that  no 
oxidation  or  reduction  occurs. 

Oxidizing  and  reducing  properties  of  oxides  and  hydroxides.  Since 
most  of  these  reactions  are  reversible,  we  may  always  expect  to  find 
an  oxide  showing  the  reactions  of  a  hydroxide  in  the  presence  of 

377 


3T8  GENERAL  CHEMISTRY 

water,  and  the  hydroxide  acting  as  an  oxide,  especially  at  high  tem- 
peratures. Consequently,  if  the  oxide  is  a  good  oxidizing  agent,  as  is 
true  of  nitrogen  pentoxide  and  sulfur  trioxide,  we  should  expect  the 
corresponding  hydroxide  (or  acid)  to  have  much  the  same  oxidizing 
properties,  and  we  have  found  this  to  be  true  of  both  nitric  and 
sulfuric  acids.  If  the  oxide  is  a  reducing  agent,  such  as  phosphorus 
trioxide,  the  hydroxide  should  have  similar  properties.  In  accord- 
ance with  this  principle  we  find  that  the  trioxide  readily  oxidizes  to 
the  pentoxide:  P.O.  +  O.-P.O.  /  ^  ' 

and  that  the  hydroxide  undergoes  an  analogous  oxidation : 
2P(OH)3  +  02=2PO(OH)3 

Dehydration  in  stages.  As  a  rule,  the  greater  the  number  of  hy- 
droxyl  groups  in  a  molecule  the  more  difficult  it  appears  to  be  for 
the  normal  hydroxide  to  exist,  and  the  tendency  to  lose  water  and 
pass  into  the  Oxide  increases.  This  transition  is  usually  incomplete  at 
ordinary  temperatures.  A  partial  dehydration  takes  place,  leading  to 
compounds  which  are  at  once  hydroxides  and  oxides,  as  indicated  in 
the  equations: 

Bi(OH)3  T=^=  BiO(OH)  +  H2O 

Si(OH)4  +=±  SiO(OH)2  +  H2O 

N(OH)5  HF±  NO(OH)3  +  H2O  +=+  NO2OH  +  2  H2O 

Acid  and  basic  conduct  as  related  to  valence.  When  the  valence  of 
an  element  is  either  1  or  2,  its  hydroxide  is  nearly  sure  to  act  as  a 
base,  though  there  are  some  exceptions.  This  means,  in  terms  of  the 
ionization  theory,  that  ionization  occurs  in  such  a  way  as  to  produce 
hydroxyl  ions.  When  the  valence  is  above  4,  the  hydroxides,  as  well 
as  the  partially  dehydrated  forms,  are  acids,  the  ionization  resulting  in 
hydrogen  ions.  With  a  valence  of  either  3  or  4  (and  sometimes  of 
only  2)  we  may  have  either  bases  or  acids.  Such  a  compound  may 
even  be  amphoteric  (p.  372),  acting  either  as  a  base  or  as  an  acid, 
according  to  the  reagents  with  which  it  is  in  contact.  A  table  will 
serve  to  illustrate  each  class. 

From  this  table  it  is  evident  that  oxygen  bases,  as  well  as  oxygen 
acids,  are  hydroxyl  compounds,  the  difference  in  valence  in  some  way 
bringing  about  a  difference  in  ionization,  and  this  in  turn  resulting  in 
a  wide  difference, in  chemical  conduct  in  the  two  classes. 


THE  HYDROXIDES   AND   THEIR  REACTIONS 


379 


CLASSES  OF  HYDROXIDES 


VALENCE 

BASE 

AMPHOTERIC 

ACID 

1 

NaOH 

2 

Ca(OH)2 

3 

Bi(OH)3 

Sb(OH)3 

P(OH)3 

3 

BiO(OH) 

SbO(OH) 

PO(OH) 

4 

Zr(OH)4 

Si(OH)4 

4 

ZrO(OH)2 

SiO(OH)2 

5 

PO(OH)3 

5 

P02(OH) 

6 

S02(OH)2 

7 

C1O3(OH) 

Action  of  elementary  substances  on  acids  and  bases.  So  far  in  the 
study  of  acids  and  bases  we  have  been  chiefly  interested  in  their 
action  upon  each  other,  that  is,  in  neutralization.  It  will  be  instructive 
to  recall,  in  a  systematic  way,  their  action  upon  elementary  substances. 

1.  Action  of  metals  upon  acids.    The  action  of  metals  upon  oxygen 
acids  may  take  place  in  either  of  two  general  ways :  If  the  acid  is  in 
rather  dilute  solution,  and  the  metal  is  above  hydrogen  in  the  electro- 
motive series,  the  metal  takes  the  place  of  hydrogen  as  a  positive  ion, 
and  hydrogen  gas  is  evolved : 

Zn  +  H2S04  =  ZnS04  +  H2 

If  the  metal  is  below  hydrogen  in  the  electromotive  series,  dilute  acids 
are  without  action  upon  it,  but  concentrated  acids,  when  they  act  at  all, 
act  as  oxidizing  agents.  A  second  reaction  frequently  accompanies  the 
oxidation,  as  is  illustrated  in  the  following  example : 

Primary :         Cu  +  H2SO4  =  CuO  +  H2SO3  — »•  H2O  +  SO2 
Secondary:  CuO  +  H2SO4  =  CuSO4  +  H2O 

2.  Action  of  nonmetals  upon  acids.    Nonmetals  do  not  replace  the 
hydrogen  of  acids,  as  they  never  form  simple  positive  ions.    Most  of 
them  are  readily  oxidizable,  however,  and  with  acids  of  marked  oxidiz- 
ing capacity  they  are  converted  into  oxides  or  hydroxides : 

C  +  2  H2S04  =  C02  +  2  H<>S03 
6  P  +  10  HNO8  =  3  P2O5  4-  5H2O  +  10  NO 


3.  Action  of  metals  upon  bases.   As  a  rule,  metals  do  not  have  a 
marked  action  upon  bases.    A  few  act  upon  the  strongest  bases  in 


380  GENERAL  CHEMISTRY 

such  a  way  as  to  displace  the  hydrogen  of  the  base.  Among  these  are 
zinc,  aluminium,  and  tin  : 

2  KOH  +  Zn  =  Zn(OK)2  +  H2 
6  KOH  +  2  Al  =  2  A1(OK)3  +  3  H2 

4.  Action  of  nonmetals  upon  bases.  The  action  of  nonmetals  upon 
strong  bases  is  of  much  interest,  since  a  number  of  important  com- 
pounds are  prepared  by  such  reactions.  The  steps  in  these  reactions 
are  open  to  some  doubt,  but  the  following  method  of  representing 
them  is  in  accord  with  the  facts,  and  at  the  same  time  shows  the  simi- 
larity of  the  reactions  with  a  variety  of  nonmetals. 

The  action  of  chlorine,  bromine,  and  iodine  is  closely  parallel,  and 
may  be  represented  by  the  equations  for  chlorine.  The  halogen  first 
sets  up  an  equilibrium  with  water  : 

HOH  +  C12  +=±  HC1  +  HC10 
The  acids  so  formed  are  then  neutralized  by  the  base  : 

HC1  +  HC1O  +  2  KOH  =  KC1  +  KC1O  +  2  H2O 

Finally,  if  the  solution  is  warm,  the  hypochlorite  undergoes  a  charac- 
teristic change  :  3  KC1Q  =  KC1O  +  2  KC1 

o 

The  complete  reaction  is  represented  in  the  equation 

6  KOH  +  3  C12  =  5  KC1  +  KC1O3  +  3  H2O 

With  sulfur  the  reactions  may  be  regarded  as  analogous,  although 
the  fact  that  sulfur  has  valences  different  from  chlorine  makes  the 
equations  different.  The  sulfur  sets  up  an  equilibrium  with  water  : 

2  HOH  +  S  +=±  S(OH)2  +  H2S 
These  acids  are  then  neutralized  by  the  base  : 

S(OH)2+  H2S  +  4  KOH  =  S(OK)2  +  K2S  +  4  H2O 

The  compound  S(OK)2  (which  is  not  known  as  an  individual  com- 
pound) then  undergoes  a  change  corresponding  to  that  of  the  hypo- 
chlorite, forming  the  sulfite  and  sulfide  : 


Finally,  both  the  sulfite  and  the  sulfide  combine  with  additional  sulfur, 
forming  thiosulfate  and  pentasulfide  : 


K2S08  +  S  =  K2S208 
K2S+4S  =  K2SE 


THE  HYDROXIDES  AND  THEIR,  REACTIONS  381 

The  complete  reaction  may  be  expressed  in  one  equation,  thus : 
6  KOH  +  128  =  K2S203  +  2  K2S5  +  3  H2O 

In  the  case  of  phosphorus  the  reaction  is  of  the  same  character, 
though  the  steps  are  not  as  clear.  The  phosphorus  acts  upon  both 
water  and  the  base,  forming  a  hypophosphite  and  phosphine : 

3  KOH  +  3  H20  +  P4  =  3  KH2PO2  +  PH8 

The  phosphine,  having  no  acid  properties,  does  not  act  with  the  base 
as  do  the  hydrides  of  the  other  nonmetals  just  considered. 

In  general,  the  action  of  a  nonmetal  upon  a  base  leads  to  the  for- 
mation of  a  hydride  of  the  nonmetal  and  a  salt  of  some  one  of  its 
oxygen  acids.  If  the  hydride  can  act  as  an  acid,  it  is  neutralized  by 
the  base  as  a  secondary  reaction. 


CHAPTER  XXVIII 

THE  METALS 

Definition.  The  elements  which  remain  to  be  considered  are  col- 
lectively called  metals,  and  the  term  at  once  suggests  a  familiar  class 
of  substances,  since  many  of  them,  such  as  gold,  silver,  copper,  and 
iron,  have  been  known  in  the  elementary  form  from  very  early  ages. 
In  their  chemical  conduct  the  metals  are  characterized  by  the  fact  that 
their  hydroxides,  at  least  those  of  lower  valence,  are  bases,  and  on  this 
account  they  are  usually  referred  to  in  chemical  literature  as  the  base- 
forming  elements.  This  is  a  more  satisfactory  term  than  metal,  since 
not  all  of  these  elements  possess  the  properties  which  are  usually  asso- 
ciated with  the  latter  term.  When  compounds  of  the  metals  are  dis- 
solved in  water,  ionization  usually  takes  place  in  such  a  way  that  the 
metal  becomes  the  cation,  while  the  remainder  of  the  compound  acts 
as  the  anion.  From  this  standpoint  the  metals  may  be  denned  as  those 
elements  which  are  capable  of  forming  simple  cations  when  their 
compounds  are  dissolved  in  water. 

Neither  of  these  definitions  is  entirely  satisfactory.  The  latter  one  depends 
upon  the  purely  theoretical  conception  of  ionization.  It  also  classes  hydrogen 
with  the  metals,  although  in  its  physical  properties  it  has  little  in  common  with 
them.  Some  of  the  metals  do  not  have  a  well-defined  valence  lower  than  3, 
and,  as  we  have  seen  in  the  preceding  chapter,  the  hydroxide  of  a  trivalent  ele- 
ment is  frequently  amphoteric  in  character.  The  majority  of  the  metals  form 
more  than  one  hydroxide,  the  one  of  higher  valence  usually  acting  as  an  acid ; 
so  not  all  of  the  hydroxides  of  the  metals  are  bases.  If  we  leave  organic  com- 
pounds out  of  consideration,  however,  it  is  only  among  the  metallic  elements 
that  we  find  hydroxides  which  are  always  bases,  and  for  this  reason  the  term 
base-forming  element  is  fairly  satisfactory  as  a  definition  of  a  metal. 

The  properties  of  the  metals.  There  are  a  number  of  properties  which, 
in  a  greater  or  less  degree,  are  characteristic  of  the  metals.  They 
reflect  light  brilliantly  from  a  polished  surface,  or  have  a  high  luster. 
With  few  exceptions,  notably  gold  and  copper,  they  have  no  individ- 
ually distinctive  colors,  but  are  all  silvery  in  appearance.  Most  of  them 
are  malleable,  or  capable  of  being  hammered  out  into  thin  foil.  They 
can  also  be  drawn  into  wire,  which  property  is  designated  as  ductility. 

382 


THE  METALS  383 

They  have  high  conductivity  for  electricity  and  heat,  and  for  the  most 
part  they  have  a  greater  density  than  the  acid-forming  elements. 

All  of  these  properties  are  greatly  modified  by  the  presence  of  even 
small  percentages  of  impurities,  as  well  as  by  the  mechanical  treatment 
to  which  the  metal  has  been  subjected.  When  cooled  from  the  liquid 
state  under  the  proper  conditions,  the  metals  are  all  highly  crystalline, 
for  the  most  part  crystallizing  in  forms  which  can  be  referred  to  the 
regular  system  of  axes.  In  such  a  state  they  are  apt  to  be  brittle  and 
to  have  little  toughness,  or  tenacity.  Under  other  conditions  of  prepa- 
ration, especially  if  hammered  or  rolled  while  cooling  (which  process 
is  called  annealing),  they  have  no  very  obvious  crystalline  form  and 
become  very  much  more  tenacious.  These  facts  show  that  it  is  not 
possible  to  assign  exact  numerical  constants  to  some  of  the  properties 
of  the  metals.  In  such  cases  the  values  given  are  merely  averages. 

Occurrence  of  the  metals  in  nature.  A  number  of  the  metals  are 
found  in  nature  in  the  uncombined,  or  elementary,  condition,  and  in 
this  case  are  said  to  occur  native.  Among  these  are  gold,  platinum, 
copper,  bismuth,  and  in  general  all  those  which  stand  low  in  the 
electromotive  series.  As  a  rule  the  metals  are  found  combined  with 
acid-forming  elements  in  the  form  of  oxides,  hydroxides,  and  salts  of 
various  acids.  The  most  abundant  of  these  salts  are  the  silicates,  car- 
bonates, sulfides,  and  sulfates.  All  such  natural  substances,  whether 
they  contain  a  metal  or  not,  are  called  minerals.  Those  minerals 
which  are  of  value  for  the  extraction  of  useful  substances,  or  find 
application  in  their  manufacture,  are  called  ores.  Most  of  these  ores 
contain  metals. 

Extraction  of  metals  from  ores ;  metallurgy.  The  art  of  extracting 
metals  from  then:  ores  is  called  metallurgy.  The  metallurgy  of  each 
metal  presents  an  individual  problem  depending  upon  the  chemical 
character  of  the  metal,  the  nature  of  the  ores  available  for  its  produc- 
tion, and  its  physical  properties,  especially  its  melting  point.  The 
problem  is  partly  a  physical  one  and  partly  chemical.  In  order  to 
obtain  the  metal  in  the  form  of  large  masses  and  in  a  state  approach- 
ing purity,  it  is  usually  necessary  to  prepare  it  at  a  temperature  above 
its  melting  point,  and  draw  it  off  from  the  furnace  in  liquid  state.  It 
is  also  desirable  that  any  earthy  impurity  entering  the  furnace  along 
with  the  ore  should  be  converted  into  a  liquid  which  can  be  easily 
removed.  .To  secure  this  end,  materials  are  mixed  with  the  ore,  which 
will  react  with  its  impurities  and  form  a  liquid  product.  The  materials 


384  GENERAL  CHEMISTRY 

so  added  are  called  fluxes,  while  the  liquid  produced  is  called  slag. 
The  latter  usually  consists  of  a  mixture  of  silicates  and  closely  re- 
sembles glass  in  character.  The  slag  also  acts  as  a  liquid  medium  in 
which  the  small  drops  of  melted  metal  can  run  together  into  larger 
masses,  and  forms  a  covering  over  the  collected  metal,  thus  protecting 
it  from  oxidation. 

While  the  details  in  each  case  vary  considerably,  there  are  a  few 
definite  principles  employed  in  metallurgy  which  cover  the  great 
majority  of  cases.  Among  these  are  the  following: 

1.  Reduction  of  an  oxide  by  carbon.  Many  of  the  metals  occur  in 
nature  in  the  form  of  oxides.  When  these  are  heated  with  carbon 
they  are  nearly  all  reduced,  as  illustrated  in  the  equations 

2  CuO  +  C  =  2  Cu  +  CO2 
Fe  O  +  3C  =  2Fe  + 


The  carbon  is  oxidized  either  to  the  dioxide  or  to  the  monoxide, 
according  to  the  temperature  which  it  is  necessary  to  employ. 

Many  ores  other  than  the  oxides  can  be  changed  into  oxides  by 
some  preliminary  treatment.  Since  some  of  the  changes  involved 
depend  upon  oxidation,  the  'process  usually  consists  in  heating  the  ore 
with  free  access  of  air,  and  is  called  roasting.  By  this  means  carbon- 
ates, hydroxides,  and  many  sulfides  are  converted  into  oxides,  as 
shown  in  the  following  typical  equations: 

4  FeC03  +  02  =  2  Fe2O3  +  4  CO2 
2Fe(OH)3=Fe203+3H20 

4FeS2  +  1102=2Fe203+8S02 

2.  Reduction  of  an  oxide  by  aluminium.  Not  all  oxides  can  be  econom- 
ically reduced  by  carbon,  however,  and  in  some  cases  in  which  the 
reduction  can  be  effected,  the  metal  combines  with  the  excess  of  car- 
bon to  form  a  carbide,  and  so  is  not  obtained  in  the  pure  metallic 
state.  In  such  cases  aluminium  may  sometimes  be  used  to  advantage 
in  place  of  carbon.  For  example,  the  element  chromium  is  prepared 
in  this  way:  Cr2O8  +  2  Al  =  A12O8  +  2  Cr 

This  method  has  come  into  use  as  a  result  of  the  cheap  production  of 
aluminium,  and  is  known  as  the  Goldschmidt  method,  since  it  was 
developed  by  the  German  chemist  of  that  name.  The  method  may  be 
illustrated  in  the  following  way  : 


THE  METALS  385 

• 

Preparation  of  chromium  by  the  Goldschmidt  method.  A  mixture  of  chromium 
oxide  and  aluminium  powder  is  placed  in  a  Hessian  crucible  (.4,  Fig.  132),  and  on 
top  of  it  is  placed  a  small  heap  (.5)  of  a  mixture  of  sodium  peroxide  and  alumin- 
ium, into  which  is  stuck  a  piece  of  magnesium  ribbon  C.  Powdered  fluor  spar 
(Z>)  is  placed  around  the  sodium  peroxide,  after  which 
the  crucible  is  set  on  a  pan  of  sand  and  the  magne- 
sium ribbon  is  ignited.  When  the  flame  reaches  the 
sodium  peroxide  mixture,  combustion  of  the  alumin- 
ium begins  with  almost  explosive  violence,  so  that 
great  care  must  be  taken  in  the  experiment.  The 
heat  of  this  combustion  starts  the  reaction  in  the 
chromium  oxide  mixture,  and  the  oxide  is  reduced 
to  metallic  chromium,  which  collects  in  the  bottom 
of  the  crucible. 

3.  Reduction  of  sulfides  by  a  metal.    Just  as  FIG 

oxygen  may  be  removed  from  an  oxide  by 

aluminium,  so  sulfur  may  sometimes  be  removed  from  a  sulfide  by 
the  same  metal  or  by  iron.  For  example,  antimony  and  lead  are 
sometimes  desulfurized  in  this  way: 

Sb  S  +  3  Fe  =  3  FeS  +  2  Sb 


•2 


4.  Electrical  methods.  Electrical  energy  is  employed  in  two  very 
different  ways  in  metallurgical  operations.  In  many  cases  it  serves 
merely  as  a  source  of  heat,  the  reduction  being  brought  about  by  car- 
bon. The  electric  furnace  has  the  advantage  of  producing  a  very  high 
temperature,  and  is  under  easy  control.  Since  the  carbon  serves  merely 
as  the  reducing  agent,  and  not  at  the  same  time  as  fuel,  its  quantity 
can  be  regulated  to  meet  the  requirements  of  the  chemical  reaction 
which  it  is  desired  to  secure.  A  typical  furnace  is  illustrated  by  the 
one  used  in  the  manufacture  of  carborundum  (p.  349). 

In  other  cases  the  electrical  energy  is  employed  to  bring  about  the 
decomposition  of  a  compound  and  the  liberation  of  the  metal  without 
the  assistance  of  a  chemical  reducing  agent.  Such  processes  are  electro- 
lytic in  character  and  always  take  place  in  a  liquid  medium.  When 
the  metal  to  be  produced  has  no  action  upon  water,  the  electrolysis 
may  be  conducted  in  aqueous  solution.  The  metal  is  then  deposited 
directly  in  solid  form  and  at  ordinary  temperatures,  and  there  is  no 
consumption  of  energy  in  maintaining  the  temperature.  In  other  cases, 
especially  those  in  which  the  metal  acts  upon  water,  some^  suitable 
salt  of  the  metal  is  melted,  and  the  resulting  electrolyte  is  subjected 
to  electrolysis.  Almost  any  melted  salt  would  serve,  but  many  are 


386  GENERAL  CHEMISTRY 

decomposed  by  heat,  or  melt  only  at  a  temperature  too  high  to  be 
economical.  In  practical  work  the  chlorides  are  usually  employed. 

Naturally  the  industries  involving  the  use  of  large  electrical  cur- 
rents tend  to  develop  at  localities  where  water  power  is  abundant, 
as  at  Niagara  Falls. 

Preparation  of  compounds  of  the  metals.  A  great  many  methods  are 
employed  in  the  preparation  of  the  compounds  of  the  metals.  These 
compounds,  which  include  oxides,  hydroxides,  and  salts,  are  very 
numerous,  and  each  has  its  own  peculiarities,  which  must  be  taken 
into  account  in  devising  means  for  preparing  it.  In  many  cases  some 
rather  unusual  method  is  employed,  owing  to  the  character  of  the 
minerals  available  in  nature  or  to  the  accumulation  of  a  cheap  by- 
product in  some  other  industry.  Naturally  the  methods  employed  on 
a  small  scale  in  the  laboratory  are  likely  to  differ  from  those  used  in 
the  industries,  where  economy  is  the  first  requirement.  There  are, 
however,  some  general  principles  which  underlie  the  great  majority 
of  these  methods,  and  it  will  save  needless  repetition  to  bring  them 
into  review  at  this  point. 

1.  Direct  union  of  two  elements.  Very  many  binary  compounds  may 
be  prepared  by  heating  the  metal  with  the   appropriate  nonmetal. 
Among  these  are  oxides,  sulfides,  and  halides.    The  product  in  such 
cases  is,  of  course,  anhydrous,  and  this  method  finds  wide  application 
when  it  is  the  anhydrous  rather  than  the  hydrated  compound  which 
is  wanted.    The  method  is  most  frequently  employed  in  the  prepara- 
tion of  anhydrous  halides,  as,  for  example,  aluminium  chloride : 

2  Al  +  3  Cl  =  2  A1C1 

&  o 

2.  Treatment  of  a  metal,  or  its  oxide  or  hydroxide,  with  an  acid.    Since 
most  of  the  metals  are  produced  commercially  in  a  high  degree  of  purity, 
the  metals  themselves  are  often  the  most  convenient  starting  point  for 
the  preparation  of  their  compounds  on  a  small  scale  in  the  laboratory. 
For  example,  the  salts  of  zinc   and  copper  are  frequently  made  in 
this  way:  Zn  +  2  HC1  =  ZnCl2  +  H2 

Cu  +  2  H2SO4  =  CuSO4  +  SO2  +  2  H2O 

In  the  industries  the  hydroxide  or  oxide  is  more  likely  to  be  employed, 
since  it  is  usually  found  in  nature  or  is  of  easy  preparation.  For  exam- 
ple, calcium  salts  are  often  prepared  from  lime  (CaO),  as  in  the  case 
of  the  nitrate  :  ' 


THE  METALS  387 

3.  Decomposition  of  compounds.    The  decomposition  of  compounds, 
either  by  heat  alone  or  in  connection  with  a  reducing  agent,  fre- 
quently   leads    to   the    formation    of    simpler    ones.    For    example, 
nitrates,  carbonates,  and  hydroxides,  when  heated  sufficiently,  usually 
yield  oxides :        2  Cu(NO3)2  -  2  CuO  +  4  NO2  +  O2 

CaCO3  =  CaO  +  CO2 
2A1(OH)3  =  A1203  +  3H20 

When  heated  with  carbon  the  various  salts  of  oxygen  acids  are 
usually  reduced  and  yield  a  binary  compound  of  the  metal  with  the 
acid-forming  element.  For  example,  sulfates  yield  sulfides  and  phos- 
phates yield  phosphides,  as  shown  in  the  equations 

BaSO4  +  2  C  =  BaS  +  2  CO2 
FePO4  +  2  C  =  FeP  +  2  CO2 

4.  Displacement  of  a  volatile  acid.    When   a  nonvolatile    acid  acts 
upon  a  salt  of  a  volatile  acid,  the  latter  is  displaced  in  accordance 
with  the  general  principles  of  equilibrium,  provided  the  volatile  acid 
is  sparingly  soluble  in  any  liquid  which  may  be  present.    For  example, 
the  reaction  of  sulfuric  acid  with  sodium  chloride  goes  on  to  conclu- 
sion if  the  sulfuric  acid  employed  is  quite  concentrated,  for  hydrogen 
chloride  is  very  sparingly  soluble  in  this  liquid: 

NaCl  +  H2S04  =  HC1  +  HNaSO4 

In  dilute  solution  the  reaction  results  in  an  equilibrium,  for  in  this 
case  the  hydrogen  chloride  is  freely  soluble  in  the  solution. 

5.  Methods  based  upon  precipitation.    The  formation  of  a  precipitate 
when  two  electrolytes  are  brought  together  in  solution  takes  place  in 
accordance   with   the    principles   of    ionic   equilibrium   developed    in 
Chapter  XIII.    In  general  it  may  be  said  that  when  two  salts,  or  a 
strong  acid  and  a  salt,  are  brought  together  in  solution,  double  decom- 
position takes  place,  with  the  formation  of  a  precipitate  if  the  union 
of  any  pair  of  ions  produces  an  insoluble  salt.    When  a  salt  is  treated 
with  a  weak  acid,  however,  no  precipitate   is  formed,  even  though 
double  decomposition  would  result  in  the  formation  of  an  insoluble 
salt.    It  can  also  be  shown  that  the  salts  of  strong  acids,  such  as  sul- 
fates or  chlorides,  when  insoluble  in  water,  are  also  insoluble  in  dilute 
acids.    Salts  of  weak  acids,  such  as  carbonates,  sulfites,  and  sulfides, 
though  insoluble  in  water,  are  soluble  in  stronger  acids.    These  prin- 
ciples are  of  constant  application  in  the  preparation  of  compounds. 


388  GENERAL  CHEMISTRY 

The  theory  of  precipitation.  A  somewhat  more  extended  discussion  of  the  prin- 
ciples just  stated  is  desirable.  It  will  be  recalled  that  all  normal  salts  (with  a 
very  few  exceptions)  are  freely  ionized  in  solution,  while  acids  differ  much  among 
themselves  in  this  respect.  This  suggests  three  general  cases  for  consideration. 

1.  Double  decomposition  between  two  salts.  As  an  example  let  the  two  salts  be  copper 
sulfate  (CuSO4)  and  barium  chloride  (BaCl2).  It  would  be  expected  that  when 
these  are  brought  together  in  solution,  all  the  salts  represented  in  the  following 
equilibrium  would  result  : 

CuSO4  +  BaCl2:«=±BaSO4  +  CuCl2 

Each  of  these  would  be  largely  ionized  and  in  equilibrium  with  its  constituent  ions, 
among  these  equilibria  being  the  following  : 

SO,-    or    ^"!  =  if  (D 


The  equation  may  also  be  written  in  the  more  convenient  form  : 

[Ba++]  x  [SO4~  ]  =  &'[BaSO4]  (2) 

When  the  equilibrium  produces  enough  barium  sulfate  to  saturate  the  solution, 
the  concentration  of  the  salt  reaches  a  fixed  value,  and  at  the  same  time  the  con- 
centration of  the  unionized  portion  [BaSO4]  reaches  a  definite  value  which  we 
may  designate  by  k".  Equation  (2)  then  becomes 

[Ba++]  x  [SO4—  ]  =  k'k"  =  K  (3) 

In  this  equation  K  is  the  product  of  the  ionization  constant  ¥  into  the  solubility 
constant  k"  (the  concentration  of  the  unionized  salt).  The  constant  K  is  called  the 
solubility  product.  We  reach  the  general  conclusion,  therefore,  that  the  product  of  the 
concentrations  of  the  ions  of  a  salt  cannot  exceed  the  solubility  product  of  the  salt. 
Of  the  two  factors  which  constitute  the  solubility  product,  kf  is  always  a  rather 
large  number,  usually  ranging  from  1  to  0.1,  since  nearly  all  salts  are  freely  ionized 
in  dilute  solution.  The  second  factor  k"  is  very  small  in  the  case  of  insoluble 
salts.  In  such  cases  the  solubility  product  does  not  differ  greatly  from  the  value 
of  k"  itself.  For  example,  barium  sulfate  is  soluble  only  to  the  extent  of  about 
2.3  mg.  per  liter,  which  is  a  molar  concentration  of  about  0.00001.  Consequently, 
the  product  [Ba++]  x  [SO4  —  ]  cannot  exceed  the  approximate  value  0.00001,  and 
when  the  ions  [Ba++]  and  [SO4  —  ]  are  brought  together  in  greater  concentrations, 
they  will  unite  and  precipitate  as  the  solid  sulfate  until  the  solubility  product  is 
no  longer  exceeded.  In  general,  therefore,  when  'two  salts  are  brought  together  in 
solution,  a  double  decomposition  takes  place  if  the  product  of  the  concentrations 
of  any  two  ions  exceeds  the  solubility  product  of  the  salt  which  is  formed  by  their 
union.  Since  all  salts  have  at  least  some  solubility,  it  is  evident  that  precipita- 
tion is  never  entirely  complete. 

2.  Double  decomposition  between  a  salt  and  a  strong  acid.  With  strong  acids  (freely 
ionized)  the  case  is  quite  similar  with  the  one  presented  by  two  normal  salts.  For 
example,  in  the  reaction  between  barium  chloride  and  sulfuric  acid  we  have,  as 

before,  the  solubility  product:    rT>  ++-, 

'  """ 


__  -,  _  v 
—  J\. 

There  is  about  as  much  of  the  ion  [SO4  —  ]  available  from  the  ionization  of  sul- 
furic acid  as  from  that  of  a  soluble  sulfate.  The  conditions  are  therefore  essen- 
tially the  same,  and  the  precipitation  is  as  complete  as  in  the  former  case. 


THE  METALS  389 

On  the  other  hand,  if  we  treat  solid  barium  sulfate  with  a  solution  of 
hydrochloric  acid,  it  does  not  dissolve  according  to  the  equation 

BaSO4  +  2  HC1  =  BaCl2  +  H2SO4 

for  this  would  at  once  bring  into  the  solution  such  a  concentration  of  the  ions 
[Ba++]  and  [SO4~  ]  as  would  exceed  the  solubility  product  [Ba++]  x  [SO4~ ]  =  A', 
when  precipitation  to  form  barium  sulfate  would  take  place  once  more. 

3.  Double  decomposition  between  a  salt  and  a  weak  acid.  When  a  salt  of  a  strong 
acid  is  treated  with  a  weak  acid,  the  case  is  quite  different.  For  example,  let  us 
suppose  that  barium  chloride  (BaCl2)  is  treated  with  carbonic  acid  (H2CO3). 
Since  barium  carbonate  is  nearly  insoluble  in  water,,  it  might  be  thought  that 
the  reaction  represented  in  the  following  equation  would  occur : 

BaCl2  +  H2C03  =  BaC03  +  2  HC1 

This  reaction  does  not  take  place  to  any  appreciable  extent,  however,  for  the 
reason  that  the  product  [Ba++]  x  [CO3 — ]  does  not  reach  the  solubility  product 
of  barium  carbonate,  owing  to  the  very  small  concentration  of  the  ion  CO3 — 
furnished  by  the  weak  carbonic  acid. 

On  the  other  hand,  when  barium  carbonate  is  exposed  to  the  action  of  strong 
acids,  such  as  hydrochloric  acid,  it  undergoes  double  decomposition,  for  it  is 
soluble  to  a  slight  extent,  and  the  dissolved  portion  freely  ionizes  into  the  ions 
Ba++  and  CO3 — .  The  latter  at  once  set  up  an  equilibrium  with  the  hydrogen 
ions  from  hydrochloric  acid: 

[H+]  x  [H+]  x  [C08~] 


[H2C03] 


=  k 


The  ionization  constant  k  is  so  small  that  nearly  all  of  the  ions  CO3 —  are  con- 
verted into  carbonic  acid  (H2CO3),  and  the  concentration  of  the  portion  remain- 
ing as  ions  is  so  small  that  the  product  [Ba++]  x  [CO3 — ]  never  reaches  the 
value  of  the  solubility  product  of  barium  carbonate.  Consequently,  the  'carbon- 
ate continues  to  dissolve  in  the  acid. 

6.  Fusion  methods.  When  compounds  are  melted  together,  it  sometimes  happens 
that  double  decomposition  occurs,  which  would  not  take  place  if  the  same  salts 
were  brought  together  in  water.  In  such  fusion  reactions  one  of  the  melted  salts 
may  be  regarded  as  the  solvent,  and  it  is  to  be  expected  that  the  solubility  of 
the  other  in  this  will  be  different  from  its  solubility  in  water.  For  example, 
barium  sulfate  is  practically  insoluble  in  water,  and  consequently  it  is  not  greatly 
affected  by  aqueous  solutions  of  sodium  carbonate.  In  melted  sodium  carbonate, 
on  the  contrary,  it  is  readily  soluble,  while  barium  carbonate  is  insoluble.  As  a 
consequence,  when  it  is  melted  together  with  sodium  carbonate,  the  following 
double  decomposition  takes  place  : 

BaSO4  +  Na2CO3  =  BaCO3  +  Na2SO4 

When  the  mixture  is  cooled  and  warmed  with  water,  all  the  compounds  except 
barium  carbonate  dissolve,  and  this  may  then  be  filtered  off  and  converted  into 
any  desired  salt.  This  forms  a  convenient  method  of  passing  from  an  insoluble 
salt  of  a  strong  acid  to  a  similar  salt  of  a  weak  acid.  It  will  be  recalled  that 
natural  silicates  can  be  decomposed  in  the  same  way. 


390  GENERAL  CHEMISTRY 

Insoluble  compounds.  From  the  foregoing  discussion  it  will  at  once 
be  evident  that  a  knowledge  of  the  solubility  of  compounds  is  an 
important  part  of  a  training  in  chemistry.  With  this  knowledge  it  is 
possible,  in  many  cases,  to  predict  the  course  of  a  reaction  and  to 
devise  ways  in  which  to  prepare  desired  compounds.  For  precise  infor- 
mation a  dictionary  of  solubilities  must  be  consulted,  but  it  is  possible 
to  make  a  few  general  statements,  covering  the  most  familiar  classes  of 
salts,  which  will  be  of  much  assistance  in  understanding  the  reactions 
of  the  chapters  which  follow.  These  statements  apply  only  to  normal 
salts  and  do  not  include  the  salts  of  rare  elements.  Acid  salts  are  apt 
to  be  more  soluble  than  normal  salts,  and  basic  salts  less  so. 

1.  Hydroxides.    All  hydroxides  are  insoluble  except  those  of  ammo- 
nium, sodium,  potassium,  calcium,  strontium,  and  barium. 

2.  Nitrates.    All  nitrates  are  soluble. 

3.  Chlorides.    All  chlorides  are  soluble  except  silver  and  mercurous 
chlorides.    (Lead  chloride  is  very  sparingly  soluble.) 

4.  Sulfates.   All  sulfates  are  soluble  except  those  of  barium,  stron- 
tium, and  lead.    (Sulfates  of  silver  and  calcium  are  only  moderately 
soluble.) 

5.  Sulfides.    All  sulfides  are  insoluble  except  those  of  ammonium, 
sodium,  and  potassium.    The  sulfides  of  calcium,  strontium,  barium, 
and  magnesium  are  insoluble  in  water,  but  are  changed  by  hydrolysis 
into  acid  sulfides  which  are  soluble.    On  this  account  they  cannot  be 
prepared  by  precipitation. 

6.  Carbonates,  sulfites,  phosphates,  and  silicates.   All  of  these  normal 
salts  are  insoluble  except  those  of  ammonium,  sodium,  and  potassium. 


CHAPTER  XXIX 

THE  ALKALI  METALS 


METAL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

BOILING 
POINT 

FIRST  PREPARED 

Lithium  (Li)  .     . 

6.94 

0.534 

186° 

1400°  + 

Arfvedson,  1817 

Sodium  (Na)  .     . 

23.00 

0.971 

97° 

877° 

Davy,  1807 

Potassium  (K)     . 

39.10 

0.862 

62.5° 

758° 

Davy,  1807 

Rubidium  (Rb)   . 

85.45 

1.532 

38.5° 

696° 

Bunsen,  1861 

Caesium  (Cs)  . 

13"2.81 

1.87 

26.4° 

670° 

Bunsen,  1861 

Characteristics  of  the  family.  The  elements  listed  in  the  above  table 
constitute  a  family  in  Group  I  of  the  periodic  table.  They  are  called 
the  alkali  metals  for  the  reason  that  the  most  familiar  members  of  the 
family,  namely,  sodium  and  potassium,  are  constituents  of  compounds 
that  have  long  been  known  as  alkalies.  Before  taking  up  the  discus- 
sion of  each  element  separately,  it  is  advisable  to  discuss  briefly  the 
family  as  a  whole. 

1.  Occurrence.  While  none  of  these  metals  occur  free  in  nature,  their 
compounds  are  widely  distributed,  being  found  in  sea  and  mineral 
waters,  in  salt  beds,  and  in  many  rocks.    Sodium  and  potassium  are, 
however,  the  only  ones  that  occur  in  abundance. 

2.  Preparation.    The  metals  are  most  readily  prepared  by  the  elec- 
trolysis of  their  fused  hydroxides  or  chlorides.    They  may  also  be 
prepared  by  the  reduction  of  their  oxides,  hydroxides,  or  carbonates. 

3.  Properties.    They  are  soft  metals,  easily  molded  by  the  fingers. 
They  have  low  melting  points  and  small  densities,  as  shown  in  the 
table.    Their  densities  (sodium  excepted)  are  in  the  same  order  as 
their  atomic  weights,  while  their  melting  and  boiling  points  are  in  the 
reverse  order.    The  pure  metals  have  a  silvery  luster  but  tarnish  at 
once  when  exposed  to  the  air,  because  of  the  formation  of  a  film  of 
oxide  upon  their  surface ;  hence  they  are  generally  preserved  in  some 
liquid,  such  as  kerosene,  which  contains  no  oxygen.    They  stand  at  the 
head  of  the  electromotive  series  of  the  metals  (p.  158)  and  in  general 
are  very  active  elements.    They  decompose  water  rapidly,  forming 

391 


392  GENERAL  CHEMISTRY 

hydroxides  and  liberating  hydrogen  in  accordance  with  the  following 
equation,  in  which  M  represents  any  one  of  these  metals : 

2  M  +  2  H2O  =  2  MOH  +  H2 

4.  Compounds.  The  alkali  metals  act  as  univalent  elements  in  the 
formation  of  compounds.  Their  hydroxides  (MOH)  are  white  solids 
and  are  very  soluble  in  water.  In  dilute  aqueous  solutions  these 
hydroxides  are  largely  ionized  and  to  about  the  same  extent,  forming 
the  ions  M+  and  OH~  ;  hence  their  solutions  are  strongly  basic.  With 
few  exceptions  the  salts  of  the  alkali  metals  are  white  solids,  and  unless 
otherwise  stated,  it  will  be  so  understood  in  the  description  of  the 
individual  compounds.  With  the  exception  of  lithium  these  metals 
form  very  few  insoluble  compounds,  so  that  it  is  difficult  to  prepare 
their  compounds  by  precipitation.  The  compounds  of  sodium  and 
potassium  are  so  similar  in  properties  that  for  most  purposes  they 
can  be  used  interchangeably.  Those  of  sodium  are  cheaper  than  the 
corresponding  ones  of  potassium,  and  so  are  more  largely  used. 

Only  sodium  and  potassium  will  be  described  in  detail,  since  the 
other  members  of  the  family  are  of  relatively  little  importance. 

LITHIUM 

The  element.  Lithium  was  discovered  in  1817  by  Arfvedson,  a 
student  of  Berzelius,  although  he  did  not  succeed  in  isolating  the  metal 
itself.  This  was  first  accomplished  by  Bunsen  in  1855.  Its  com- 
pounds are  widely  but  sparingly  distributed,  being  found  in  nearly 
all  igneous  rocks,  from  which  it  finds  its  way  into  the  soil.  Certain 
plants,  such  as  the  sugar  beet  and  tobacco,  absorb  small  quantities  of 
lithium  compounds  from  the  soil,  and  when  such  plants  are  burned, 
the  lithium  remains  in  the  ash  in  the  form  of  the  carbonate.  Lith- 
ium is  also  present  in  most  mineral  waters  (lithia  waters),  although 
usually  in  traces  only.  Some  of  the  most  important  of  the  lithium 
minerals  are  lepidolite,  found  in  California,  and  spodumene  and 
amblygonite,  found  in  South  Dakota.  These  contain  from  4  to  10 
per  cent  of  lithium. 

Lithium  may  be  prepared  by  the  electrolysis  of  its  fused  chloride 
or  of  a  solution  of  the  chloride  in  some  solvent,  such  as  pyridine,  which 
is  not  acted  upon  by  the  metal.  It  is  the  lightest  of  all  the  elements 
which  are  solid  at  ordinary  temperature,  having  a  density  of  only 
0.534.  It  resembles  the  other  alkali  metals  in  properties.  Lithium 


THE  ALKALI  METALS  393 

unites  with  nitrogen  even  at  ordinary  temperatures  and  is  therefore 
sometimes  used  in  place  of  magnesium  in  the  preparation  of  argon 
(p.  108).  The  metal  is  very  expensive  and  has  no  commercial  uses. 

Compounds  of  lithium.  The  element  forms  many  compounds,  but 
only  a  very  few  are  of  any  commercial  importance.  When  heated  in 
a  Bunsen  flame,  most  of  these  compounds  volatilize  and  impart  to  the 
flame  a  crimson  color.  Even  when  the  quantity  present  is  so  small  that 
the  flame  is  not  visibly  colored,  the  presence  of  the  element  may  be 
recognized  by  viewing  the  flame  through  the  spectroscope.  Lithium 
compounds  are  prepared  by  dissolving  lithium  minerals  in  acids  and 
precipitating  lithium  from  the  solution  in  the  form  of  the  carbonate. 
This  can  be  changed  into  other  salts  by  the  action  of  the  appropriate 
acid.  Some  of  the  most  important  of  the  lithium  compounds  are  the 
following : 

Lithium  chloride  (LiCl)  is  prepared  directly  from  the  lithium  min- 
erals and,  like  the  carbonate,  is  used  in  the  preparation  of  other 
lithium  compounds.  Lithium  bromide  (LiBr)  is  used  in  medicine. 
Lithium  carbonate  (Li2CO3)  is  also  used  in  medicine  and  in  the  con- 
struction of  the  Edison  storage  battery.  It  is  but  slightly  soluble  in 
water  and  is  one  of  the  few  compounds  whose  solubility  decreases 
with  rise  in  temperature  — 100  g.  of  water  at  20°  dissolving  1.33  g.  of 
the  carbonate  and  at  100°,  0.73  g.  Lithium  phosphate  (Li3PO4),  like  the 
carbonate,  is  but  slightly  soluble  in  water.  The  sparing  solubility  of 
the  carbonate  and  phosphate  is  noteworthy,  since,  with  these  exceptions, 
all  the  common  compounds  of  the  alkali  metals  are  readily  soluble. 

SODIUM 

History.  The  isolation  of  sodium  dates  back  to  the  year  1807.  At 
that  time  the  compounds  now  known  as  sodium  hydroxide  and  potas- 
sium hydroxide  were  well  known  and  were  called  "  fixed  alkalies," 
but  they  were  regarded  as  elementary  in  character.  In  1807  Sir  Hum- 
phry Davy,  while  studying  the  effect  of  the  electric  current  upon 
various  substances,  succeeded  in  decomposing  these  fixed  alkalies 
and  thus  obtained  metallic  sodium  and  potassium. 

Davy  announced  his  discovery  in  a  letter  to  a  friend,  as  follows :  w  I  have 
decomposed  and  recomposed  the  fixed  alkalies  and  discovered  their  bases  to  be 
two  new  inflammable  substances  very  like  metals ;  but  one  of  them  lighter  than 
ether  and  infinitely  combustible.  So  that  there  are  two  bodies  decomposed  and 
two  new  elementary  bodies  found." 


394 


GENERAL  CHEMISTRY 


Occurrence.  Sodium  occurs  as  a  silicate  in  many  igneous  rocks, 
especially  in  the  feldspar  known  as  albite.  By  the  disintegration  of 
these  rocks  compounds  of  sodium  pass  into  the  soil,  from  which  they 
are  taken  up  by  plants,  although  not  in  such  large  quantities  as  are 
the  potassium  compounds.  The  most  familiar  compound  of  sodium 
is  the  chloride  which  occurs  in  all  salt  and  mineral  waters  and  also 
forms  large  deposits  in  various  parts  of  the  world.  Other  sodium 
compounds  found  in  nature  are  the  nitrate  (Chile  saltpeter),  the 
carbonate,  the  sulfate,  and  the  borate  (borax). 

Preparation.  For  many  years  the  most  economical  method  known 
for  preparing  sodium  consisted  in  the  reduction  of  its  carbonate : 

Na2COs  4-20=2  Na +  3  CO 

At  present  it  is  all  prepared  by  the  electrolysis  of  either  the  fused 
hydroxide  or  chloride.  It  is  evident  that  all  water  must  be  excluded 
in  the  process ;  otherwise  the  sodium  liberated  will  react  with  the 
water  and  form  sodium  hydroxide. 

Castner's  process.  At  Niagara  Falls  sodium  is  prepared  by  the  electrolysis  of 
fused  sodium  hydroxide  by  a  process  devised  by  Castner.  The  apparatus  consists 
of  a  cylindrical  iron  vessel  A,  A  (Fig.  133),  through  the  bottom  of  which  extends 
an  iron  rod  B,  which  serves  as  the  cathode.  The  iron  anodes  C,  C,  several  in 

number,  are  suspended  around  the  cathode  but  are 
kept  from  touching  it  by  a  cylinder  of  iron  gauze 
D,  which  is  fastened  to  the  vessel  E.  The  lower 
part  of  the  vessel  A,  A  is  filled  with  molten  sodium 
hydroxide,  which,  on  cooling,  holds  the  cathode  in 
position.  The  heat  generated  by  the  current  is  ordi- 
narily sufficient  to  keep  the  hydroxide  in  the  upper 
portion  of  the  vessel  fused ;  however,  the  apparatus 
is  supplied  with  a  row  of  gas  burners  G,  G,  which 
may  be  utilized  if  additional  heat  is  required. 
Sodium  and  hydrogen  are  liberated  at  the  cathode 
and,  rising  to  the  surface,  collect  in  the  vessel  E. 
The  hydrogen  escapes  by  lifting  the  cover  of  the 
vessel,  while  the  sodium,  protected  from  the  air  by 
the  hydrogen,  is  skimmed  or  drawn  off  from  time 
to  time.  Oxygen  is  liberated  at  the  anodes  and 
escapes  through  the  opening  F  without  coming  in 
contact  with  either  the  sodium  or  the  hydrogen. 

Properties  and  uses.  Sodium  is  a  soft,  silver-white  metal,  slightly 
lighter  than  water.  It  melts  at  97°  and  boils  at  877°.  It  is  very 
active  chemically,  combining  readily  with  most  of  the  nonmetallic 
elements  such  as  oxygen  and  the  halogens;  It  decomposes  water 


FIG.  133 


THE  ALKALI  METALS  395 

and  reacts  with  acids,  forming  the  corresponding  salts  and  liberat- 
ing hydrogen.  It  dissolves  in  mercury,  forming  an  alloy  (sodium 
amalgam)  which  is  an  efficient  reducing  agent.  When  heated  in  the 
Bunsen  flame,  sodium,  as  well  as  most  of  its  compounds,  volatilizes 
and  imparts  a  yellow  color  to  the  flame  —  a  property  which  is  used 
as  a  test  for  the  presence  of  the  element. 

Sodium  is  used  in  the  preparation  of  sodium  cyanide  and  sodium 
peroxide  and  to  a  limited  extent  as  a  reducing  agent. 

Compounds  of  sodium.  With  the  exception  of  the  nitrate  all  the 
compounds  of  sodium  are  prepared  from  the  chloride,  since  it  is  so 
abundant  and  inexpensive.  The  processes  involved  are  often  compli- 
cated, owing  to  the  fact  that  the  compounds  of  sodium  are  all  soluble 
and  therefore  cannot  be  prepared  directly  from  the  chloride  by  pre- 
cipitation ;  moreover,  the  chloride  is  a  salt  of  a  strong  acid  and  is  not 
readily  acted  upon  by  most  other  acids.  Experiments  have  shown 
that  the  most  economical  method  of  procedure  consists  either  in  first 
changing  the  chloride  into  the  hydroxide  by  the  electrolysis  of  its 
aqueous  solution  (p.  143),  or  in  converting  it  into  the  carbonate  by 
the  methods  to  be  described.  Since  the  hydroxide  is  a  base  and  the 
carbonate  is  a  salt  of  a  very  volatile  acid,  both  are  readily  changed 
into  other  compounds. 

Sodium  hydride  (NaH).  Moissan  obtained  this  compound  in  the 
form  of  white  crystals  by  heating  sodium  and  hydrogen  at  360°. 
Since  it  is  easily  decomposed  into  its  elements,  both  of  which  have  a 
strong  affinity  for  oxygen,  it  is  a  powerful  reducing  agent.  It  reacts 
with  water  and  acids,  as  indicated  in  the  following  equations  : 

NaH  +  H2O  ^  NaOH  +  H2 


The  oxides  of  sodium.  The  metal  forms  two  oxides,  namely,  sodium 
oxide  (Na.2O)  and  sodium  peroxide  (Na2O2).  The  former  is  obtained, 
mixed  with  the  peroxide,  when  sodium  is  burned  in  a  limited  supply 
of  air.  The  peroxide  is  more  easily  obtained  pure,  and  is  of  much  the 
greater  importance,  since  it  serves  as  an  excellent  oxidizing  agent.  It 
is  a  yellowish-white  powder  prepared  by  passing  air,  freed  from  mois- 
ture and  carbon  dioxide,  through  an  iron  tube  containing  sodium 
heated  to  about  300°.  While  the  pure  compound  is  stable  toward 
heat,  in  the  presence  of  an  oxidizable  substance  it  gives  up  half  of 
its  oxygen,:  NaO=NaO 


•2~2          '"a    • 


396 


GENERAL   CHEMISTRY 


It  is  readily  acted  upon  by  water  and  acids : 

Na202  +  2  H20  =  2  NaOH  +  H2O2 
Na2O2  +  2  HC1  =  2  NaCl  +  H2O2 

The  hydrogen  peroxide  formed  in  the  reactions  decomposes  into  water 
and  oxygen  unless  the  temperature  is  kept  low.  It  will  be  recalled 
that  the  reaction  with  water  serves  as  one  of  the  methods  used  for 
preparing  oxygen  (p.  20). 

Sodium  hydroxide  (caustic  soda)  (NaOH).  This  compound  is  pre- 
pared on  a  large  scale  by  two  general  processes. 

1.  Action  of  calcium  hydroxide  upon  sodium  carbonate.  This  process 
consists  in  treating  calcium  hydroxide  suspended  in  water  with  sodium 
carbonate.  Calcium  carbonate,  being  insoluble,  is  precipitated : 

Na2CO8  +  Ca(OH)2  =  CaCO3  +  2  NaOH 

The  resulting  sodium  hydroxide  is  obtained  by  filtering  off  the  calcium 
carbonate  and  evaporating  the  filtrate  to  dryness. 

While  this  is  an  old  process  it  still  remains  the  chief  one  for  the  production 
of  the  hydroxide.  Manufacturers  of  sodium  carbonate  often  utilize  a  portion  of 
their  product  in  the  preparation  of  the  hydroxide,  so  that  the  manufacture  of 
these  two  compounds  is  often  carried  out  in  the  same  plant. 

2.  Electrolytic  methods.  By  the  newer  method 
sodium  hydroxide  is  obtained  by  the  electrolysis 
of  sodium  chloride.  The  products  of  the  electrol- 
ysis are  sodium  hydroxide,  hydrogen,  and  chlo- 
rine (p.  143).  The  chief  difficulty  in  this  process 
is  to  prevent  the  chlorine  and  the  hydroxide  from 
acting  upon  each  other.  This  is  usually  done  by 
separating,  by  means  of  a  porous  diaphragm,  the 
anode  and  cathode  compartments  of  the  cell  in 
which  the  electrolysis  is  effected.  A  number  of 
different  cells  have  been  devised  for  carrying  out 
the  process  ;  at  present  one  of  the  most  successful 
of  these  is  that  devised  by  Townsend  and  known 
by  his  name. 


> 


n 


FIG.  134 


Commercial  preparation  by  the  Townsend  cell.  A  section 
of  this  cell  is  shown  in  Fig.  134.  The  anode  compartment 
is  formed  by  the  diaphragm  A,  A,  a  nonconducting  bottom  B,  and  a  lid  C.  The 
diaphragm  is  made  of  asbestos  cloth  painted  over  with  a  mixture  of  iron  oxide 
and  asbestos  fiber.  Through  the  lid  C  extends  the  graphite  anode  D.  The 


THE  ALKALI  METALS  397 

diaphragm  is  set  firmly  against  the  perforated  iron  cathode  plate  E,  E,  which  is 
in  turn  held  in  place  by  the  iron  sides  F,  F,  the  space  between  the  plate  and  the 
iron  sides  forming  the  cathode  compartment.  The  anode  compartment  is  par- 
tially filled  with  saturated  salt  solution  G,  and  the  cathode  compartment  with 
kerosene  H.  Since  the  level  of  the  salt  solution  is  above  that  of  the  kerosene, 
the  solution  slowly  penetrates  the  diaphragm,  and  some  of  the  salt,  coming  in 
contact  with  the  cathode,  is  changed  into  the  hydroxide.  The  resulting  solution 
of  the  chloride  and  hydroxide  enters  the  anode  compartment  and,  being  heavier 
than  the  kerosene,  sinks  to  the  bottom  and  is  drawn  off  through  the  side  tubes. 
The  chloride,  being  much  less  soluble  than  the  hydroxide,  is  separated  by  partial 
evaporation  of  the  solution.  The  hydrogen  and  the  chlorine  that  are  set  free  are 
led  off  through  tubes,  and  the  chlorine  is  used  in  the  preparation  of  bleaching  pow- 
der. The  hydrogen  is  sometimes  used  in  preparing  hydrochloric  acid  (p.  253). 

Sodium  hydroxide  is  a  crystalline,  brittle  solid  which  rapidly  absorbs 
water  and  carbon  dioxide  from  the  air,  being  changed  thereby  into  the 
carbonate.  As  the  name,  caustic  soda,  indicates,  it  is  a  corrosive  sub- 
stance and  has  a  disintegrating  action  upon  most  animal  and  vegetable 
tissues.  It  is  a  strong  base  and  is  used  in  many  chemical  industries 
such  as  the  manufacture  of  soap  and  paper. 

Sodium  chloride  (common  salt)  (NaCl).  Sodium  chloride  is  very  widely 
distributed  in  nature.  Thick  strata,  evidently  deposited  by  the  evapora- 
tion of  salt  water,  are  found  in  many  places.  In  the  United  States  the 
most  important  localities  for  salt  are  New  York,  Michigan,  Ohio,  and 
Kansas.  Sometimes  the  salt  is  mined,  especially  if  it  is  in  the  pure  form 
called  rock  salt.  More  frequently  a  strong  brine  is  pumped  from  deep 
wells  sunk  into  the  salt  deposit,  and  is  then  evaporated  in  large  pans 
until  the  salt  crystallizes  out.  The  crystals  are  in  the  form  of  small 
cubes  and  contain  no  water  of  crystallization  ;  some  water  is,  however, 
held  in  cavities  in  the  crystals  and  causes  the  salt  to  decrepitate  when 
heated.  It  melts  at  801°  and  above  this  temperature  begins  to  volatilize. 

Salt  is  used  in  the  preparation  of  nearly  all  substances  containing 
either  sodium  or  chlorine.  These  include  many  products  of  the  highest 
importance  to  civilization,  such  as  soap,  glass,  hydrochloric  acid,  soda, 
and  bleaching  powder.  Enormous  quantities  of  salt  are  therefore  pro- 
duced each  year.  Small  quantities  are  essential  to  animal  life.  Pure 
salt  does  not  absorb  moisture ;  the  fact  that  ordinary  salt  becomes 
moist  in  air  is  due  to  the  presence  in  it  of  certain  deliquescent  com- 
pounds, especially  calcium  and  magnesium  chlorides. 

Preparation  of  pure  sodium  chloride.  The  salt  found  in  nature  is,  with  few  ex- 
ceptions, more  or  less  impure.  Pure  sodium  chloride  may  be  prepared  by  the 
action  of  hydrochloric  acid  upon  sodium  carbonate;  or  the  impurities  may  be 


398  GENERAL  CHEMISTRY 

removed  by  passing  hydrogen  chloride  into  a  saturated  solution  of  the  impure 
salt,  pure  sodium  chloride  being  thereby  precipitated. 

This  precipitation  of  the  chloride  is  easily  understood  from  the  following 
considerations  :  in  the  salt  solution  the  molecules  of  salt  are  in  equilibrium  with 
the  ions  Na+  and  Cl~  : 


XT   ~, 

NaCl  < 

If  the  solution  is  saturated,  we  have  the  following  equation  (p.  388)  : 

[Na+]  x  [C1-]  =  K 

If  hydrogen  chloride  is  passed  into  such  a  solution,  the  number  of  chlorine  ions 
is  increased.  The  number  of  sodium  ions  must  therefore  correspondingly  de- 
crease, in  order  that  the  value  of  K  may  remain  constant.  This  decrease  is  brought 
about  by  some  of  the  sodium  ions  combining  with  chlorine  ions  to  form  molecules 
of  sodium  chloride.  Since  the  solution  is  already  saturated  with  this  compound, 
the  excess  is  precipitated. 

Sodium  bromide  (NaBr)  ;  sodium  iodide  (Nal).  These  compounds 
somewhat  closely  resemble  sodium  chloride  in  their  physical  properties. 
They  can  be  prepared  by  the  action  of  bromine  and  iodine  respectively 
upon  a  solution  of  sodium  hydroxide  (p.  380).  They  are  used  to  a 
limited  extent  in  photography. 

Sulfides  of  sodium.  Sodium  forms  both  the  acid  and  the  normal  salt 
with  hydrosulfuric  acid,  namely,  NaHS  and  Na2S.  The  former  is  pre- 
pared by  saturating  a  solution  of  sodium  hydroxide  with  hydrogen  sul- 
fide,  and  the  latter  by  reducing  sodium  sulf  ate  with  carbon.  The  normal 
salt,  when  dissolved  in  water,  is  largely  hydrolyzed,  giving  a  strong 
alkaline  solution.  When  sulfur  is  heated  with  a  solution  of  sodium 
hydroxide,  sulfides  which  contain  more  than  one  atom  of  sulfur  to  the 
molecule  are  formed,  and  are  known  collectively  as  the  polysulfides  of 
sodium.  They  are  difficult  to  purify,  so  that  their  composition  has 
not  been  definitely  determined. 

Sodium  sulf  ate  (Na2S04).  Large  quantities  of  this  compound  are 
prepared  for  use  in  the  manufacture  of  glass  and  sodium  carbonate. 
It  is  obtained  by  the  action  of  sulf  uric  acid  upon  sodium  chloride: 

H2SO4  +  2  NaCl  =  Na2SO4  +  2  HC1 

It  will  be  recalled  that  this  reaction  serves  as  the  principal  source  of 
hydrochloric  acid  (p.  254  ).  Sodium  sulfate  is  also  prepared  by  the 
action  of  sodium  chloride  upon  magnesium  sulfate,  the  latter  being 
obtained  in  large  quantities  as  a  by-product  in  the  manufacture  of 
potassium  chloride  : 

MgSO4  +  2  NaCl  =  Na2SO4  +  MgCl2 


THE  ALKALI  METALS  399 

The  sodium  sulfate  formed  in  this  reaction  is  the  least  soluble  of  the 
compounds  represented  in  the  equation,  and  separates  when  hot  satu- 
rated solutions  of  magnesium  sulfate  and  sodium  chloride  are  mixed. 

Under  ordinary  conditions  sodium  sulfate  crystallizes  from  water 
in  the  form  of  a  decahydrate  (Na2SO4  -10  H2O).  This  is  known  as 
G-laubers  salt,  from  the  alchemist  Glauber,  who  lived  about  the  middle 
of  the  seventeenth  century,  and  who  first  used  the  compound  in  medi- 
cine. Considerable  deposits  of  this  hydrate  are  found  in  nature.  If 
a  supersaturated  solution  of  the  sulfate  is  cooled,  the  heptahydrate 
(Na2SO4  •  7  H2O)  separates.  The  transition  point  between  the  deca- 
hydrate and  the  anhydrous  salt  is  32.38°,  and  this  point  is  so  definite 
that  it  has  been  suggested  by  Richards  and  Wells  as  a  suitable  fixed 
temperature  for  use  in  the  calibration  of  thermometers. 

By  the  action  of  sulfuric  acid,  sodium  sulfate  is  converted  into  the 
acid  sulfate  NaHSO4,  commonly  known  as  bisulfate  of  sodium. 

Sulfites'of  sodium.  The  acid  sulfite  NaHSO3,  often  called  sodium 
bisulfite,  is  formed  by  saturating  a  solution  of  sodium  carbonate  with 
sulfur  dioxide.  Sulfurous  acid  is  first  formed  by  the  union  of  the 
dioxide  with  water,  and  this  decomposes  the  carbonate: 

Na2C03  +  2  H2S03  =  2  NaHS08  +  C02  +  H20 

The  normal  sulfite  Na2SO3  is  prepared  by  adding  sodium  carbonate 
to  a  saturated  solution  of  the  acid  sulfite  in  the  proportion  indicated 
in  the  following  equation  : 

2  NaHS03  +  Na2C03  =  2  Na2SO3  -h  H2O  +  CO2 

Both  of  the  sulfites  readily  absorb  oxygen,  forming  the  corresponding 
sulfates  ;  they  are  therefore  reducing  agents.  They  are  used  to  some 
extent  as  bleaching  agents  and  as  preservatives. 

Sodium  thiosulfate  (Na2S203).  This  salt  is  made  by  adding  sulfur  to 
a  solution  of  sodium  sulfite  and  warming  the  mixture  : 


Upon  evaporation  the  resulting  solution  yields  the  pentahydrate 
Na2S2O3  •  5  H2O,  which  is  commonly  known  as  hyposulfite  of  soda, 
or  simply  hypo.  It  is  also  prepared  from  the  calcium  sulfide  obtained 
as  a  by-product  in  the  manufacture  of  sodium  carbonate.  Upon  exposure 
to  air  the  calcium  sulfide  is  oxidized  to  calcium  thiosulfate,  which  is 
then  treated  with  sodium  carbonate  : 

CaS203  +  NasC08  =  Na,SaO8  +  CaCO3 


400  GENERAL  CHEMISTRY 

The  insoluble  calcium  carbonate  is  filtered  off,  and  the  sodium  thio- 
sulfate is  obtained  by  evaporating  the  filtrate. 

Sodium  thiosulfate  is  very  soluble  in  water,  in   which  it  readily 
forms  supersaturated  solutions.    Acids  react  with  it,  liberating  free 

sulfur :  Na2S203  +  2  HC1  =  2  NaCl  +  H2S2O3 


Dilute  solutions  exposed  to  the  air  undergo  a  similar  change  due  to 
the  action  of  carbonic  acid.  The  salt  reacts  with  free  iodine,  forming 
sodium  iodide  and  sodium  tetrathionate  : 

2  Na2S203  +  I2  =  Na2S406  +  2  Nal 

This  reaction  is  used  for  determining  the  quantity  of  free  iodine  present  in  a 
solution.  The  liquid  is  first  colored  blue  by  the  addition  of  a  few  drops  of  starch 
solution  (p.  265).  A  solution  of  sodium  thiosulfate  of  known  strength  is  then 
slowly  run  in  from  a  burette  until  the  blue  color  just  disappears,  thus  indicating 
that  all  the  iodine  has  entered  into  combination.  From  the  amount  of  sodium 
thiosulfate  added  one  can  easily  calculate  the  quantity  of  iodine  present. 

The  salt  is  used  very  largely  in  photography  as  a  solvent  for  silver 
salts  (see  photography)  and  as  an  "  antichlor  "  for  removing  any  chlorine 
remaining  in  substances  bleached  with  this  element. 

It  must  be  noted  that  the  name  Jiyposulfite  as  applied  to  this  compound  is 
a  misnomer.  If  the  general  system  of  naming  salts  were  followed,  the  term 
sodium  hyposulfite  would  represent  the  sodium  salt  of  hyposulfurous  acid. 

Sodium  carbonate  (soda  ash)  (Na2C03).  This  very  important  com- 
pound occurs  in  nature  to  a  limited  extent  in  certain  arid  regions. 
Many  seaweeds  are  rich  in  sodium  compounds  absorbed  from  the 
water.  When  these  are  burned,  sodium  carbonate,  along  with  other 
sodium  salts,  remains  in  the  ashes ;  hence  the  name  soda  ash.  At  present 
it  is  all  made  from  sodium  chloride  by  one  of  two  general  methods. 

1.  Leblanc  process.  This  older  process,  no  longer  in  use  in  the 
United  States  but  still  used  in  Europe,  involves  several  distinct  reac- 
tions, the  most  important  ones  being  represented  in  the  following 
equations : 

O  ^u^  £.  /i  1       |      TT    O  f\ ^f        C*  f\          I 

Na2SO4  +  20  =  NaJS  +  2  CO2 
Na0S  +  CaCO0  =  CaS  +  NaOO, 


2 


In  a  manufacturing  plant  the  last  two  reactions  take  place  in  one  process. 
Sodium  sulf  ate,  coal,  and  powdered  limestone  are  heated  together  to  a  rather  high 
temperature.  The  coal  reduces  the  sulfate  to  sulfide,  which  in  turn  reacts  upon 


THE  ALKALI  METALS  401 

the  calcium  carbonate.  Some  limestone  is  decomposed  by  the  heat,  forming 
calcium  oxide.  When  treated  with  water  the  calcium  oxide  is  changed  into  hy- 
droxide, and  this  prevents  the  insoluble  calcium  sulfide  from  being  hydrolyzed 
into  the  soluble  calcium  acid  sulfide  Ca(HS)2,  which  would  react  with  the  sodium 
carbonate  to  form  the  insoluble  calcium  carbonate. 

The  crude  product  of  the  process  is  a  hard  black  cake  called  black  ash.  When 
this  mass  is  digested  with  water,  the  sodium  carbonate  passes  into  solution.  The 
pure  carbonate  is  obtained  by  evaporation  of  this  solution,  crystallizing  from  it 
in  crystals  of  the  formula  Na2CO3  •  10  H2O.  Since  over  60  per  cent  of  this  salt 
is  water  of  crystallization,  the  crystals  are  usually  heated  until  it  is  driven  off. 
The  product  is  called  calcined  soda  or  soda  ash. 

2.  Solvay  process.  This  more  modern  process  depends  upon  the 
reactions  represented  in  the  equations 


NaCl  +  NH4HC08  =  NaHCO3  +  NH4C1  (1) 

2  NaHCO8  =  Na2CO3  +  H2O  4-  CO2  .  (2) 

When  concentrated  solutions  of  sodium  chloride  and  of  ammonium 
hydrogen  carbonate  are  brought  together,  the  sparingly  soluble  sodium 
hydrogen  carbonate  is  precipitated,  as  represented  in  equation  (1). 
This  is  converted  into  the  normal  carbonate  by  heating,  the  reaction 
being  represented  in  equation  (2). 

In  the  Solvay  process  a  very  concentrated  solution  of  salt  is  first  saturated 
with  ammonia  gas,  and  a  current  of  carbon  dioxide  is  then  conducted  into  the 
solution.  In  this  way  ammonium  hydrogen  carbonate  is  formed  : 

NH3  +  H20  +  C02  -  NH4HC03 

This  enters  into  double  decomposition  with  the  salt,  as  shown  in  equation  (1) 
under  the  Solvay  process.  After  the  sodium  hydrogen  carbonate  has  been  precipi- 
tated, the  mother  liquors  containing  ammonium  chloride  are  treated  with  lime  : 

2  NH4C1  -I-  CaO  =  CaCl2  +  2  NH8  +  H2O 
The  lime  is  obtained  by  burning  limestone  : 

CaCO3  =  CaO  +  CO2 

The  ammonia  and  carbon  dioxide  evolved  in  these  reactions  are  used  in  the 
preparation  of  an  additional  quantity  of  ammonium  hydrogen  carbonate.  There 
is,  therefore,  no  loss  of  ammonia  ;  the  only  materials  permanently  used  up  are 
salt  and  lime,  while  the  only  by-product  is  calcium  chloride. 

Historical.  In  former  times  sodium  carbonate  was  obtained  principally  from 
the  ashes  of  certain  plants.  During  the  French  Revolution  this  supply  was  cut 
off,  and  in  behalf  of  the  French  government  Leblanc  made  a  study  of  methods 
of  preparing  the  carbonate  directly  from  salt.  As  a  result  he  devised  the  method 
which  bears  his  name,  and  which  was  used  exclusively  for  many  years.  It  has 
been  replaced  to  a  large  extent  by  the  Solvay  process,  which  has  the  advantage 


402  GENERAL   CHEMISTRY 

that  salt  and  lime  are  inexpensive,  and  that  the  ammonium  hydrogen  carbonate 
used  can  be  regenerated  from  the  products  formed  in  the  process.  Much  expense 
is  also  saved  in  fuel,  and  the  sodium  hydrogen  carbonate,  which  is  the  first 
product  of  the  process,  has  itself  many  commercial  uses.  The  Leblanc  process 
is  still  used  to  a  limited  extent,  however,  since  the  hydrochloric  acid  generated 
in  the  process  is  a  valuable  by-product. 

Properties  and  uses  of  sodium  carbonate.  The  aqueous  solution  of 
the  salt  is  basic  in  character  (p.  225).  With  water  the  salt  forms  a 
number  of  different  hydrates,  the  most  common  of  which  is  the  deca- 
hydrate  Na2CO3'10  H2O.  By  cooling  a  hot  solution  of  the  salt  this 
hydrate  is  obtained  in  the  form  of  large,  clear,  monoclinic  crystals  and 
is  often  known  as  washing  soda,  or  sal  soda.  It  effloresces  on  exposure 
to  dry  air,  changing  into  the  monohydrate.  The  transition  point  of 
the  decahydrate  into  the  heptahydrate  Na2CO8  •  7  H2O  is  32°,  and  of 
this  into  the  monohydrate,  35.37°.  The  anhydrous  salt  melts  at  853° 
and  along  with  potassium  carbonate  is  used  for  decomposing  silicates 
(p.  346).  Mere  mention  of  the  fact  that  sodium  carbonate  is  used  in 
the  manufacture  of  glass,  soap,  and  many  chemical  reagents  will  indi- 
cate its  importance  in  the  industries.  Enormous  quantities  of  it  are 
manufactured  for  these  various  uses. 

Sodium    hydrogen    carbonate    (bicarbonate  of  soda)    (baking  soda) 
(NaHCOg).   This  salt  is  prepared  either  by  the  Solvay  process,  as  al- 
ready explained,  or  by  passing  carbon  dioxide  into  saturated  solutions 
of  sodium  carbonate:  J^ ^fl    ("1  CL£Jj-f  H^    "'  Y]f\  ?C  <b^ 
Na2C03  +  H20  +  C02  =  2  NaHCO3 

The  bicarbonate,  being  but  sparingly  soluble,  precipitates.  When 
heated,  the  bicarbonate  changes  into  the  carbonate,  with  liberation  of 
carbon  dioxide.  The  salt  is  used  as  an  aerating  agent  in  baking.  For 
this  purpose  it  must  be  mixed  with  some  substance,  such  as  sour  milk 
or  cream  of  tartar  (p.  306),  which  slowly  reacts  with  the  carbonate, 
liberating  carbon  dioxide. 

Cream  of  tartar  baking  powders.  Cream  of  tartar  baking  powders  consist  of  a 
mixture  of  cream  of  tartar,  bicarbonate  of  soda,  and  some  starch  or  flour.  When 
water  is  added  to  this  mixture,  the  cream  of  tartar  slowly  acts  upon  the  soda, 
liberating  carbon  dioxide  in  accordance  with  the  following  equation : 

KHC4H406  +  NaHC03  =  KNaG4H4O6  +  II2O  +  CO2 

The  carbon  dioxide  escapes  through  the  dough,  making  it  light  and  porous.  The 
starch  is  added  to  absorb  any  moisture  present  in  the  other  ingredients  of  the 
powder,  and  thus  to  prevent  their  interaction  until  the  powder  is  used. 


THE  ALKALI  METALS  403 

Sodium  nitrate  (Chile  saltpeter)  (NaN03).  This  substance  is  found 
in  certain  arid  regions,  where  it  has  apparently  been  formed  by  the 
decay  of  organic  substances  in  the  presence  of  air  and  sodium  salts. 
The  largest  deposits  are  in  Chile,  and  most  of  the  nitrate  of  commerce 
comes  from  that  country.  Smaller  deposits  occur  in  California  and 
Texas.  The  crude  nitrate  is  known  as  caliche.  The  commercial  salt 
is  prepared  by  treating  the  caliche  with  water,  allowing  the  insoluble 
earthy  materials  to  settle,  and  evaporating  to  crystallization  the  clear 
solution  so  obtained.  The  soluble  impurities  remain  for  the  most  part 
in  the  mother  liquors. 

Since  this  salt  is  the  only  nitrate  found  extensively  in  nature,  it  is 
the  material  from  which  other  nitrates,  as  well  as  nitric  acid,  are  pre- 
pared. It  is  used  in  enormous  quantities  in  the  manufacture  of  sulfuric 
acid  and  potassium  nitrate,  and  as  a  fertilizer. 

Sodium  cyanide  (NaNC).  This  salt  of  hydrocyanic  acid  possesses  the 
property  of  dissolving  gold  and  is  used,  along  with  potassium  cyanide, 
for  extracting  this  metal  when  it  is  scattered  in  small  quantities 
through  earthy  material.  It  is  prepared  by  heating  sodium  f errocyanide 
(NaJFeN  C  )  with  sodium: 

v          4  6      6/ 

Na FeN  C  +  2  Na  =  6  NaNC  +  Fe 

466 

It  is  also  prepared  from  sodamide  (p.  172)  by  heating  it  with  carbon : 
NaNH2  +  C  =  NaNC  +  H2 

Its  aqueou*  solution  is  strongly  alkaline  (p.  225).  Like  hydrocyanic 
acid,  it  is  extremely  poisonous. 

Sodium  phosphates.  Since  phosphorus  forms  a  number  of  acids, 
most  of  which  are  poly  basic,  one  can  readily  understand  why  so  many 
different  phosphates  of  sodium  are  known.  The  names  and  formulas 
of  the  sodium  salts  of  orthophosphoric  acid  have  already  been  given 
(p.  361),  and  only  these  compounds  will  be  discussed. 

1.  Normal  sodium  phosphate  (Na3POJ.  Although  this  te  a  normal 
salt,  its  aqueous  solution  has  a  strong  basic  reaction,  due  to  partial 
hydrolysis :  Na3PO4  +  H2O  :<=>:  Na2HPO4  +  NaOH 

It  is  prepared  by  adding  sodium  hydroxide  to  a  solution  of  disodium 
phosphate  and  evaporating  to  crystallization.  The  excess  of  sodium 
hydroxide  reverses  the  reaction  of  hydrolysis,  and  the  normal  salt 
separates  in  the  form  of  crystals  having  the  formula  NagPO4  •  12  H2O. 
The  salt  is  sometimes  used  in  laundries  for  softening  water.  It  not 


404  GENERAL   CHEMISTRY 

only  precipitates  the  calcium  and  magnesium  salts  present,  but  at  the 
same  time  leaves  the  water  slightly  basic  in  reaction. 

2.  Disodium  phosphate  (Na2HPO^).    This  is  the  most  common  of  the 
phosphates  of  sodium  and  is  generally  known  simply  as  sodium  phos- 
phate.   It  occurs  in  blood  and  urine,  and  it  was  from  these  sources  that 
the  salt  was  first  obtained.    It  is  prepared  by  the  action  of  phosphoric 
acid  upon  sodium  carbonate : 

Na2C03  +  H3P04  =  Na2HP04  +  H2O  +  CO2 

The  salt  crystallizes  from  solution  in  the  form  of  the  hydrate 
Na2HPO4  •  12H2O.  This  is  the  salt  commonly  used  when  a  soluble 
phosphate  is  needed. 

3.  Monosodium  phosphate  (NaHzPO^.    This  salt  is  prepared  by  the 
action  of  phosphoric  acid  upon  disodium  phosphate  : 

H3PO4  +  Na2HPO4  =  2  NaH2PO4 

Sodium  pyroantimonate  (Na2H2Sb207  •  6  H20).  This  salt  is  precipitated 
when  a  solution  of  potassium  pyroantimonate  is  added  to  a  concen- 
trated solution  of  a  sodium  compound.  It  is  the  least  soluble  of  all 
the  compounds  of  sodium. 

POTASSIUM 

Occurrence.  Potassium  is  a  rather  abundant  element,  being  a  con- 
stituent of  many  igneous  rocks,  especially  the  feldspars  and  micas.  Sea 
water,  as  well  as  most  mineral  waters,  contains  small  percentages  of  its 
compounds.  Very  large  deposits  of  the  chloride  and  sulfate,  asso- 
ciated with  compounds  of  calcium  and  magnesium,  occur  at  Stassfurt, 
Germany,  and  are  known  as  Stassfurt  salts.  It  is  also  found  in  small 
quantities  as  the  nitrate  (saltpeter)  and  in  many  other  forms. 

The  natural  decomposition  of  rocks  containing  potassium  gives  rise 
to  various  compounds  of  the  element  in  all  fertile  soils.  It  is  absorbed 
by  growing  plants  and  is  a  characteristic  constituent  of  land  plants, 
just  as  sodium  is  of  sea  plants.  Some  of  the  sea  plants,  however,  as, 
for  example,  the  giant  algae  of  the  California  coast,  contain  potassium 
chloride  amounting  in  some  cases  to  30  per  cent  of  their  dry  weight. 
In  the  land  plants  the  potassium  is  present  chiefly  in  the  form  of  salts 
of  organic  acids.  When  such  plants  are  burned,  the  potassium  remains  in 
the  ash  as  carbonate,  and  the  crude  carbonate  so  obtained  was  formerly 
the  chief  source  of  potassium  compounds.  At  present,  however,  they 
are  prepared  almost  entirely  from  the  salts  of  the  Stassfurt  deposits. 


THE  ALKALI  METALS 


405 


Stassfurt  salts.  These  salts,  evidently  deposited  from  sea  water  under  peculiar 
geological  conditions,  form  very  extensive  deposits  in  middle  and  north  Germany, 

the  most  noted  locality  for  working  them      

being  at  Stassfurt.  The  deposits  are  very 
thick  and  rest  upon  an  enormous  layer  of 
common  salt.  They  are  in  the  form  of 
a  series  of  strata,  each  consisting  largely 
of  a  single  mineral  salt.  Over  thirty 
different  minerals  are  present,  although 
some  in  very  small  quantities.  Fig.  135 
shows  a  cross  section  of  these  deposits. 
While  from  a  chemical  standpoint  these 
strata  are  salts,  they  are  as  solid  and 
hard  as  many  kinds  of  stone  and  are 
mined  as  stone  or  coal  would  be.  Since 
the  strata  differ  in  general  appearance, 
each  can  be  mined  separately,  and  the 
various  minerals  can  be  worked  up  by 
methods  adapted  to  each  particular  case. 
The  chief  minerals  of  commercial  impor- 
tance in  these  deposits  are  the  following : 


FIG.  135 


Sylvite     .     .     .     KC1  Kainite 

Anhydrite    .     .     CaSO4  Kieserite 

Carnallite     .     .     KC1  •  MgCl2  •  6  H2O         Schonite 


KC1-3H20 


MgS04 

MgS04-H20 
K2S04-MgS04-6H20 


Preparation  and  properties.  The  metal  can  be  obtained  by  the  general 
method  used  in  the  preparation  of  sodium.  This  process,  however,  is 
more  difficult  to  carry  out,  and  as  the  metal  has  no  particular  uses, 
but  little  of  it  is  produced.  It  has  a  density  of  0.862,  melts  at  62.5°, 
and  boils  at  758°.  It  is  very  similar  to  sodium,  differing  from  it 
mainly  in  its  greater  activity.  It  decomposes  water  violently,  the  heat 
of  the  reaction  being  sufficient  to  ignite  the  hydrogen  evolved. 

Compounds  of  potassium.  In  a  general  way  the  compounds  of  potas- 
sium are  similar  to  the  corresponding  ones  of  sodium  and  therefore 
will  not  be  discussed  in  such  detail.  Many  of  the  compounds  volatilize 
when  heated  in  a  Bunsen  flame,  and  impart  to  it  a  characteristic  violet 
color,  which  serves  to 'indicate  the  presence  of  the  element.  If  other 
compounds  which  mask  this  color  are  present  (for  example,  those  of 
sodium),  the  flame  may  be  examined  through  the  spectroscope,  the 
characteristic  spectrum  of  potassium  being  easily  recognized. 

Potassium  hydroxide  (caustic  potash)  (KOH).  This  compound  is  pre- 
pared by  the  methods  used  in  the  production  of  sodium  hydroxide. 
It  is  very  soluble  in  water,  the  solution-  being  strongly  basic.  Exposed 


406  GENERAL  CHEMISTEY 

to  the  air  it  rapidly  absorbs  water  and  is  a  good  dehydrating  agent. 
It  is  often  used  in  the  laboratory  to  remove  both  water  and  carbon 
dioxide  from  gases.  It  is  not  used  to  any  great  extent  commercially, 
being  replaced  by  the  cheaper  sodium  hydroxide. 

Potassium  halides.  Of  these  compounds  potassium  chloride  is  the 
most  familiar,  since  it  is  found  in  such  large  quantities  in  the  Stass- 
fiirt  deposits.  The  mineral  sylvite  is  nearly  pure  potassium  chloride. 
The  salt  is  obtained  not  only  from  sylvite  but  also  from  carnallite 
(KC1  •  MgCl2  •  6  H2O).  When  dissolved  in  water,  carnallite  separates 
into  its  constituent  compounds,  and  advantage  is  taken  of  this  fact  in 
the  preparation  of  potassium  chloride.  A  hot  saturated  solution  of  the 
mineral  is  first  prepared.  When  this  is  cooled,  the  potassium  chloride, 
being  less  soluble  than  the  magnesium  chloride,  crystallizes  out.  In  its 
general  properties  potassium  chloride  resembles  sodium  chloride.  It 
is  used  in  the  preparation  of  nearly  all  other  potassium  salts  and  as 
a  fertilizer.  Potassium  bromide  (KBr)  is  prepared  by  the  action  of 
bromine  upon  a  hot  solution  of  potassium  hydroxide  (p.  380). 

6  KOH  +  3  Br2  =  5  KBr  +  KBrO3  +  3  H2O 

By  heating  the  product  the  bromate  is  converted  into  the  bromide, 
with  evolution  of  oxygen,  so  that  only  the  pure  bromide  remains.  It 
is  also  prepared  commercially  by  treating  a  bromide  of  iron  (Fe3Brg) 
with  potassium  carbonate : 

Fe3Brg  +  4  K2CO8  =  Fe3O4  +  8  KBr  +  4  CO2 

Potassium  iodide  (KI)  is  prepared  by  the  same  methods  as  those  used 
in  preparing  potassium  bromide.  Both  the  iodide  and  the  bromide  are 
used  in  photography  and  in  medicine. 

Potassium  chlorate  (KC103).  This  salt  is  formed  by  the  action  of 
chlorine  upon  warm  solutions  of  potassium  hydroxide  (p.  380). 

3  C12  +  6  KOH  =  5  KC1  +  KC1O8  +  3  H2O 

It  will  be  noted,  however,  that  the  yield  is  very  small,  six  molecules 
of  the  hydroxide  giving  but  one  of  the  chlorate.  Commercially  the 
yield  is  greatly  improved  by  generating  the  chlorine  and  potassium 
hydroxide  by  the  electrolysis  of  potassium  chloride  under  such  condi- 
tions that  they  react  to  form  the  chlorate  according  to  the  above 
equation.  By  continuing  the  process  all  the  chloride  is  finally  con- 
verted into  the  chlorate. 


THE  ALKALI  METALS  407 

Another  process  used  commercially  consists  in  the  action  of  potassium  chloride 
upon  calcium  chlorate,  the  latter  compound  being  prepared  from  calcium  hydroxide 
(slaked  lime)  at  very  low  cost : 

Ca(C103)2  +  2  KC1  -  2  KC1O3  +  CaCl2 

It  will  be  noted  that  all  the  potassium  entering  into  the  reaction  is  converted 
into  the  chlorate.  The  potassium  chlorate  is  separated  from  the  accompanying 
calcium  chloride  by  evaporating  the  solution.  The  chlorate,  being  much  less 
soluble  than  the  calcium  chloride,  separates  first. 

Potassium  chlorate  melts  at  370°.  When  heated  to  a  higher  tem- 
perature, a  portion  of  the  compound  is  converted  into  potassium  chlo- 
ride and  oxygen,  while  another  portion  is  converted  into  potassium 
chloride  and  perchlorate.  When  treated  with  hydrochloric  acid,  the 
oxygen  of  the  chlorate  unites  with  the  hydrogen  of  the  acid,  thus 
liberating  chlorine.  This  mixture  may  therefore  be  used  in  place  of 
aqua  regia  as  a  solvent.  The  chief  use  of  potassium  chlorate  is  as  an 
oxidizing  agent  in  the  manufacture  of  matches,  fireworks,  and  explo- 
sives. It  is  also  used  in  medicine  and  in  the  preparation  of  oxygen. 

Potassium  bromate  (KBr03)  and  potassium  iodate  (KI03).  These  compounds  are 
made  by  methods  similar  to  those  used  in  preparing  the  chlorate.  Like  the  chlo- 
rates, they  are  strong  oxidizing  agents. 

Potassium  sulfate  (K2SOJ.  This  salt  is  formed  by  the  action  of  sul- 
furic  acid  upon  potassium  chloride.  Commercially  it  is  prepared  from 
the  Stassfurt  salts,  especially  schonite,  by  the  action  of  potassium 

K2S04  •  MgS04  +  2  KC1  -  2  K2SO4  +  MgCla 

When  the  solution  is  evaporated,  the  resulting  potassium  sulfate, 
being  much  less  soluble  than  the  magnesium  chloride,  separates  first. 
Potassium  sulfate  is  used  as  a  fertilizer  and  for  making  potassium 
aluminium  sulfate  (potash  alum).  When  heated  with  sulfuric  acid, 
it  is  converted  into  the  bisulfate  KHSO4. 

Potassium  carbonate  (K2C03).  This  compound  can  be  prepared  from 
potassium  chloride  by  the  Leblanc  process,  just  as  sodium  carbonate 
is  prepared  from  sodium  chloride.  Commercially  it  is  chiefly  prepared 
according  to  the  reactions  indicated  in  the  following  equations : 

3  MgC08  +  2  KC1  +  C02  +  H20  =  2  MgKH(CO3)2  +  MgCl2  (1) 
The  resulting  carbonate  is  decomposed  when  heated: 

2  MgKH(C03)2  =  2  MgCO,  +  K2CO3  +  CO2  +  H2O  (2) 


408  GENERAL   CHEMISTRY 

The  magnesium  carbonate  and  carbon  dioxide  formed  according  to 
equation  (2)  react  with  a  further  supply  of  potassium  chloride  accord- 
ing to  equation  (1),  and  the  process  thus  continues.  If  equations  (1) 
and  (2)  are  combined,  the  resulting  equation  is  as  follows : 

MgC03  +  2  KC1  =  MgCl2  +  K2C03 

This  shows  that  magnesium  carbonate  and  potassium  chloride  are  the 
only  substances  used  up  in  the  reaction. 

.  Potassium  carbonate  is  used  in  the  manufacture  of  glass  and,  to  a 
limited  extent,  in  the  preparation  of  other  compounds  of  potassium. 
When  carbon  dioxide  is  passed  into  a  saturated  solution  of  the 
carbonate,  potassium  bicarbonate  (KHCO3)  is  formed.  Both  the  car- 
bonate and  bicarbonate  are  very  similar  to  the  corresponding  salts 
of  sodium. 

Potassium  nitrate  (saltpeter)  (KN03).  This  compound  constituted 
one  of  the  most  important  reagents  of  the  alchemists.  It  is  formed  in 
the  decay  of  nitrogenous  organic  matter  (p.  189)  and  therefore  accumu- 
lates in  some  regions  where  the  climate  is  hot  and  dry.  At  present  it 
is  prepared  by  the  action  of  sodium  nitrate  upon  potassium  chloride : 

NaNO,  +  KC1  =  NaCl  +  KNO 

o  3 

The  sodium  nitrate  used  in  the  process  is  obtained  from  the  Chile 
niter  beds  and  the  potassium  chloride  from  the  Stassfurt  salts. 

The  reaction  depends  for  its  success  upon  the  apparently  insignificant  fact  that 
sodium  chloride  is  almost  equally  soluble  in  cold  and  in  hot  water.  All  four  com- 
pounds represented  in  the  equation  are  rather  soluble  in  cold  water,  but  in  hot 
water  sodium  chloride  is  far  less  soluble  than  the  other  three.  When  hot  satu- 
rated solutions  of  sodium  nitrate  and  potassium  chloride  are  brought  together, 
sodium  chloride  precipitates  and  can  be  filtered  off,  leaving  potassium  nitrate  in 
solution,  together  with  some  sodium  chloride.  When  cooled,  potassium  nitrate 
crystallizes  out,  leaving  small  amounts  of  the  other  salts  in  solution. 

Potassium  nitrate  dissolves  in  water  with  marked  absorption  of 
heat.  It  crystallizes  from  water  in  large  rhombic  crystals  which  melt 
at  345°.  When  heated  alone  it  gives  up  oxygen,  forming  the  nitrite. 
It  was  in  this  way  that  Scheele  first  obtained  oxygen.  It  is  an  excel- 
lent oxidizing  agent.  Its  chief  use  is  in  the  manufacture  of  gun- 
powder (p.  331)  ;  for  this  purpose  it  is  preferable  to  sodium  nitrate, 
since  the  latter  is  deliquescent,  and  powder  made  with  it,  if  exposed 
to  air,  soon  becomes  unfit  for  use.  Smaller  amounts  are  used  in  medicine' 
and  as  a  preservative  for  meat. 


THE  ALKALI  METALS  409 

Potassium  cyanide  (KNC).  Potassium  cyanide  is  very  similar  to 
sodium  cyanide  in  its  properties  and  is  prepared  by  the  same  general 
processes.  When  the  compound  known  as  potassium  ferrocyanide 
(K4Fe(NC)6)  is  heated  to  a  red  heat,  it  decomposes,  forming  potassium 
cyanide,  iron  carbide,  and  nitrogen  : 

K4Fe(NC)6  =  4  KNC  +  FeC2  +  N2 

The  yield  is  improved  and  a  purer  product  obtained  by  heating  a 
mixture  of  the  ferrocyanide  and  potassium: 

K4Fe(NC)6  +  2  K  =  6  KNC  +  Fe 

Since  sodium  is  much  cheaper  than  potassium,  it  is  often  used  in  place 
of  it,  the  product  being  a  mixture  of  the  cyanides  of  sodium  and 
potassium  :  KFe(NC)  +  2  Na  =  4  KNC  +  2  NaNC  +  Fe 


This  mixture  is  known  commercially  as  potassium  cyanide  and  serves 
for  most  of  the  purposes  of  the  pure  salt.  It  is  used  especially  in  the 
extraction  of  gold  from  earthy  materials  (p.  538).  The  cyanides  must 
be  used  with  extreme  precaution,  since  they  are  not  only  exceedingly 
poisonous  in  themselves,  but  in  contact  with  almost  any  of  the  acids 
they  evolve  the  deadly  poisonous  fumes  of  hydrocyanic  acid. 

With  appropriate  oxidizing  agents  potassium  cyanide  yields  potas- 
sium cyanate  (KNCO)  ;  with  sulfur  it  yields  potassium  sulfocyanate 
(KNCS).  The  latter  serves  as  a  very  delicate  reagent  for  the  detec- 
tion of  certain  compounds  of  iron,  since  it  reacts  with  them  to  form 
the  deep  red  sulfocyanate  of  iron  (Fe(NCS)3). 

Insoluble  compounds  of  potassium.  The  following  compounds  of  potassium  are 
but  slightly  soluble  in  water,  so  that  the  metal  may  be  precipitated  from  its 
solutions  in  these  forms  :  (1)  potassium  perchlorate  (KC1O4),  a  white,  crystal- 
line solid;  (2)  potassium  chloroplatinate  (K2PtCl6),  a  yellow,  crystalline  solid; 

(3)  potassium  sodium  cobaltinitrite  (K2NaCo(NO2)6),  a  yellow,  crystalline  solid; 

(4)  potassium  fluosilicate  (K2SiF6),  a  white  solid.    The  solubility  of  all  these 
compounds  is  greatly  decreased  by  the  addition  of  alcohol. 

RUBIDIUM  AND  CAESIUM 

These  two  elements  were  discovered  by  Bunsen  while  making  a 
spectroscopic  examination  of  the  residues  from  certain  mineral  waters. 
The  characteristic  lines  in  the  spectrum  of  the  one  are  red  in  color, 
while  those  of  the  other  are  blue  ;  hence  the  names  rubidium,  meaning 
"  dark  red,"  and  ccesium,  meaning  "  blue." 


410  GENERAL  CHEMISTRY 

Rubidium  and  caesium  are  generally  associated  with  potassium,  al- 
though present  in  very  small  quantities.  Rubidium  is  absorbed  from 
the  soil  by  certain  plants,  especially  the  sugar  beet  and  tobacco. 
Caesium  occurs  on  the  island  of  Elba  in  the  form  of  the  very  rare 
mineral  known  as  pollucite,  which  is  a  caesium  aluminium  silicate. 

The  free  metals  are  very  difficult  to  prepare ;  their  most  important 
properties  have  been  given  in  the  table  at  the  beginning  of  the  chapter. 
They  form  compounds  analogous  in  formulas  and  general  properties 
to  those  of  sodium  and  potassium.  Because  of  their  high  cost  neither 
the  metals  nor  their  compounds  have  any  commercial  uses. 

COMPOUNDS  OF  AMMONIUM 

General.  As  explained  in  Chapter  XV,  when  ammonia  is  passed 
into  water,  the  two  unite  to  form  the  base  ammonium  hydroxide,  and 
when  this  base  is  neutralized  with  acids,  ammonium  salts  are  formed. 
Since  the  ammonium  group  is  univalent,  ammonium  salts  resemble 
those  of  the  alkali  metals  in  formulas ;  they  also  resemble  the  latter 
salts  in  their  chemical  properties,  and  may  be  conveniently  described 
in  connection  with  them.  They  all  volatilize  upon  being  heated,  most 
of  them  being  decomposed  in  the  process.  When  heated  with  an 
aqueous  solution  of  sodium  hydroxide,  they  evolve  ammonia  (p.  168). 
Since  the  ammonia  can  be  easily  recognized,  the  reaction  serves  for 
the  detection  of  the  presence  of  ammonium  compounds. 

Ammonium  amalgam.  While  the  ammonium  radical  NH4  has  never  been  iso- 
lated in  the  pure  state,  it  is  easy  to  prepare  an  amalgam  which  apparently  con- 
sists of  a  solution  of  ammonium  in  mercury.  The  discovery  of  this  amalgam 
dates  back  to  1808,  when  Seebeck  noted  that  if  an  electric  current  is  passed 
through  a  solution  of  aqua  ammonia  to  which  some  mercury  has  been  added,  the 
mercury  is  greatly  increased  in  bulk  and  acquires  the  properties  of  an  amalgam 
somewhat  resembling  sodium  amalgam  in  its  general  characteristics.  Later  it 
was  found  that  a  similar  product  could  be  obtained  by  the  electrolysis  of  any 
ammonium  salt,  mercury  being  used  as  the  cathode.  The  amalgam-like  product 
soon  decomposes,  evolving  ammonia  and  hydrogen  and  leaving  the  pure  mercury- 

Occurrence.  Small  quantities  of  ammonium  compounds  are  found  in 
the  soil.  They  are  being  continually  absorbed  by  growing  plants,  but 
are  returned  to  it  again  in  the  process  of  decay.  They  are  also  found  in 
sea  water  and  in  some  volcanic  regions.  Larger  quantities  are  found 
in  the  Stassfurt  deposits.  Commercially  ammonium  compounds  are 
all  prepared  from  the  ammoniacal  liquors  produced  in  the  manufac- 
ture of  coal  gas  (p.  323). 


THE  ALKALI  METALS  411 

Ammonium  chloride  (sal  ammoniac)  (NH4C1).  This  compound  was 
known  and  used  by  the  ancients,  who  obtained  it  by  burning  animal 
excrement.  It  is  prepared  commercially  by  treating  the  ammoniacal 
liquors  of  the  gas  works  with  lime  and  passing  into  hydrochloric  acid 
the  ammonia  which  is  evolved.  By  evaporating  the  resulting  solution 
the  impure  salt  is  obtained.  This  is  purified  by  sublimation. 

The  density  of  the  vapor  obtained  by  heating  ammonium  chloride 
is  only  about  half  what  one  would  expect  if  the  vapor  consisted  of 
molecules  of  the  salt.  Experiments  have  shown  that  this  apparent 
discrepancy  is  due  to  the  fact  that  at  high  temperatures  the  salt  is 
dissociated  into  ammonia  and  hydrogen  chloride,  which,  however, 
recombine  as  the  temperature  falls: 


The  salt  is  used  in  soldering,  since  the  hydrogen  chloride  evolved  in 
the  process  removes  any  oxide  from  the  surface  of  the  metals.  It  is 
also  used  in  medicine,  in  the  preparation  of  ammonia,  in  making  dry 
cells,  and  as  a  chemical  reagent. 

Other  ammonium  halides.  These  resemble  ammonium  chloride  in  their  general 
properties.  The  iodide  KH4I  readily  absorbs  moisture  upon  exposure  to  air 
and  decomposes  to  such  an  extent  that  the  free  iodine  liberated  colors  the  salt. 
The  fluoride  (NH4)2F2  attacks  silicates  and  is  used  for  etching  upon  glass.  The 
bromide  and  iodide  are  used  in  photography. 

Ammonium  sulfides.  The  normal  sulfide  ((NH4)2S)  and  the  acid 
sulfide  ((NH4)HS)  are  formed  when  hydrogen  sulfide  and  ammonia 
are  brought  together  in  the  proper  proportions  at  temperatures  below 
zero.  They  form  colorless  crystals  which  dissociate  into  ammonia  and 
hydrogen  sulfide  as  the  temperature  rises.  In  solution  they  are  formed 
by  passing  hydrogen  sulfide  into  aqua  ammonia  : 

2  NH4OH  +  H2S  =  (NH4)2S  +  2  H2O 
NH4OH  +  H2S  =  (NH4)HS  -f  H2O 

The  normal  salt,  however,  is  almost  completely  hydrolyzed  in  solution, 
forming  the  acid  sulfide  and  ammonium  hydroxide. 

The  solution  obtained  by  passing  hydrogen  sulfide  into  aqua  ammonia, 
and  commonly  known  as  ammonium  sulfide,  is  largely  used  in  the  laboratory  as 
a  reagent  in  the  precipitation  of  certain  metals.  When  exposed  to  air  this  solu- 
tion gradually  decomposes,  the  hydrogen  sulfide  formed  in  the  hydrolysis  being 
oxidized  to  water  and  sulfur.  The  sulfur,  however,  does  not  separate  but  com- 
bines with  the  compounds  present,  forming  several  different  sulfides,  such  as 
(NH4)2S2,  (NHJ2S3,  (NH4)2S5.  The  resulting  solution  is  yellow  in  color  and  is 


412  GENEKAL  CHEMISTRY 

known  as  ammonium  polysulfide,  or  yellow  ammonium  sulfide.  It  is  used  in  the 
laboratory  as  a  solvent  for  the  sulfides  of  arsenic,  antimony,  and  tin.  It  can  be 
prepared  by  adding  sulfur  to  a  solution  of  ordinary  ammonium  sulfide.  Some  of 
the  individual  polysulfides  have  been  obtained  in  the  form  of  pure  crystals  that 
are  fairly  stable. 

Ammonium  sulfate  ((NHJ2S04).  This  is  one  of  the  cheapest  and 
most  widely  used  of  the  ammonium  salts.  It  is  prepared  by  passing 
ammonia  into  sulfuric  acid  and  is  largely  used  as  a  fertilizer.  By  the 
action  of  sulfuric  acid  it  is  changed  into  the  bisulfate  NH4HSO4 ;  this 
upon  electrolysis  yields  ammonium  persulfate  ((NH4)2S2Og),  which  is 
used  as  an  oxidizing  agent. 

Ammonium  carbonates.  When  a  mixture  of  limestone  (CaCO3) 
and  ammonium  chloride  is  heated,  there  is  formed  as  a  sublimate  a 
compound  which  is  made  up  of  ammonium  bicarbonate  and  ammo- 
nium carbamate  (p.  288)  and  has  the  formula  NH4HCO3  •  NH4CO2NH2. 
This  salt  is  known  as  commercial  ammonium  carbonate.  When  ammonia 
is  passed  into  a  concentrated  aqueous  solution  of  this  salt,  the  normal 
ammonium  carbonate  (NH4)2COg  is  formed  and,  being  but  slightly 
soluble  in  strong  ammonia  water,  separates  as  a  white,  crystalline  solid : 

NH4HC03  -  NH4C02NH2  +  HaO  +  NH3  =  2  (NH4)2CO3 

The  normal  carbonate  is  unstable,  decomposing  at  ordinary  tempera- 
tures with  evolution  of  ammonia: 

(NH4)2C03  =  NH4HC03  +.NH, 

The  bicarbonate  (NH4HCO3)  is  much  more  stable.  It  is  prepared  by 
passing  carbon  dioxide  into  aqua  ammonia.  When  heated  sufficiently 
it  is  decomposed  into  water  and  the  gases,  ammonia  and  carbon  di- 
oxide, so  that  it  is  sometimes  used  as  an  aerating  agent  in  making 
certain  forms  of  pastry.  A  solution  of  the  carbonates  is  used  in  the 
laboratory  in  reactions  requiring  a  soluble  carbonate. 

Ammonium  nitrate  (NH4N03).  This  salt  may  be  prepared  by  the 
action  of  nitric  acid  upon  either  ammonium  hydroxide  or  carbonate. 
It  is  used  in  the  preparation  of  nitrous  oxide  (p.  185)  and  as  a  con- 
stituent of  certain  explosives. 


CHAPTER  XXX 


THE  ALKALINE  EARTH  METALS 


ATOMIC 

WJ^T 

D^^ 

MELTING 

POINT^ 

MELTING  POINT 
OF  CHLORIDE 

Calcium  (Ca)            .     . 

40.1 

1.55 

780° 

780° 

Strontium  (Sr)    .     .     „ 

87.63 

2.54 

900° 

800°  + 

Barium  (Ba) 

137.37 

3.75 

850° 

960°  — 

The  family.  Calcium,  strontium,  and  barium  are  known  as  the 
alkaline  earth  metals.  Together  with  radium  they  constitute  one  of 
the  families  in  Group  II  of  the  periodic  table.  While  radium  closely 
resembles  barium,  it  is  more  convenient  to  discuss  it  in  connection 
with  uranium,  to  which  it  bears  a  peculiar  relation.  The  term  alkaline 
earths  was  originally  applied  to  the  oxides  of  these  metals  because  they 
bore  some  resemblance  both  to  the  alkalies  and  to  the  earths,  the  latter 
being  a  general  term  for  such  oxides  as  those  of  iron  and  aluminium. 
As  in  the  case  of  the  alkalies,  the  alkaline  earths  were  thought  to  be 
elementary  in  character  until  1807,  when  Davy  succeeded  in  decom- 
posing them  just  as  he  had  decomposed  the  alkalies. 

In  a  lecture  delivered  on  June  30,  1808,  before  the  English  Royal  Society, 
Davy  refers  to  his  discovery  as  follows :  "  The  evidence  for  the  composition  of 
the  alkaline  earths  is  then  of  the  same  kind  as  that  for  the  composition  of  the 
common  metallic  oxides ;  and  the  principles  of  their  decomposition  are  precisely 
similar.  .  .  .  These  new  substances  will  demand  new  names ;  and  on  the  same 
principles  as  I  have  named  the  bases  of  the  fixed  alkalies  I  shall  venture  to  de- 
nominate the  metals  from  the  alkaline  earths  barium,  strontium,  calcium.  .  .  ." 

1.  Occurrence.    Like  the  alkali  metals,  the  alkaline  earth  metals  do 
not  occur  free  in  nature.    Their  most  abundant  compounds  are  the 
carbonates  and  sulfates,  calcium  also  occurring  in  large  quantities  in 
the  form  of  the  phosphates  and  silicate. 

2.  Preparation.    The  metals  are  prepared  by  the  electrolysis  of  their 
melted  chlorides  or  hydroxides.   Calcium  is  the  most  readily  prepared. 

3.  Properties.    The  three  metals  resemble  one  another  very  closely. 
They  are  silvery  white  in  color  and  are  somewhat  harder  than  lead. 

413 


414 


GENERAL  CHEMISTRY 


Like  the  alkali  metals,  they  combine  readily  with  oxygen  and  there- 
fore tarnish  upon  exposure  to  air.  They  decompose  water  at  ordinary 
temperatures,  forming  hydroxides  and  liberating  hydrogen,  although 
not  so  readily  as  do  the  alkali  metals.  When  ignited  in  air  they  burn 
with  brilliancy,  forming  oxides  of  the  general  formula  MO,  in  which 
M  represents  any  one  of  the  metals.  These  oxides  combine  with  water 
to  form  hydroxides  of  the  general  formula  M(OH)2. 

4.  Compounds.  The  alkaline  earth  metals  act  as  divalent  elements 
in  the  formation  of  salts.  The  corresponding  salts  of  the  three  ele- 
ments are  similar  to  one  another  and  show  a  regular  gradation  in  many 
of  their  properties.  Unlike  the  alkali  metals,  their  normal  carbonates, 
phosphates,  and  silicates  are  insoluble  in  water.  Barium  sulfate  is  also 
insoluble,  while  the  sulf  ates  of  calcium  and  strontium  are  but  sparingly 
soluble.  When  volatilized  in  a  colorless  flame,  the  compounds  of  each 
of  the  three  metals  impart  a  characteristic  color  to  the  flame,  those  of 
calcium  giving  a  light  red  color,  those  of  strontium  a  deeper  red,  and 
those  of  barium  a  green  color. 

CALCIUM 

Occurrence.  Compounds  of  calcium  are  found  in  large  quantities  in 
various  regions.  The  most  abundant  of  these  compounds  is  the  car- 
bonate, which  occurs  in  many  different  forms,  such  as  marble  and 
limestone.  Other  calcium-bearing  minerals  are  the  following:  fluora- 

patite  (3  Ca3(PO4)2  •  CaF2),  chlorapa- 
tite  (3  Ca3(PO4)8  •  CaCl2),  fluor  spar 
(CaF^,  wollastonite  (CaSiO^,  gypsum 
(CaSO4  •  2  H2O),  anhydrite  (CaSO4), 
and  phosphorite  (Ca3(PO4)2). 

Preparation.  Davy  first  isolated 
calcium  by  the  electrolysis  of  the 
hydroxide.  It  is  now  prepared  by 
the  electrolysis  of  the  melted  chlo- 
ride, and  a  number  of  different  cells 
have  been  devised  for  effecting  the 
electrolysis. 

Method  of  Seward  and  von  Kugelgen.  One 
FIG.  136  form  of  cell  for  the  commercial  production 

of   calcium  is  represented   in  Fig.  136.   It 

consists  of  a  cylindrical  iron  vessel,  through  the  bottom  of  which  extends  the  iron 
cathode  A.  The  anodes  B,  B,  several  in  number,  are  distributed  about  the  sides 


THE  ALKALINE  EARTH  METALS  415 

of  the  vessel.  The  calcium  separates  at  the  cathode  in  a  molten  condition  and 
rises  in  the  form  of  globules  to  the  lower  surface  of  a  solid  stick  of  calcium  Z>, 
suspended  above  the  cathode,  as  shown  in  the  diagram.  There  it  becomes  chilled 
by  a  water-cooling  device  C,  C  and  adheres  to  the  stick  of  calcium,  which  is  slowly 
raised  as  it  increases  in  length. 

Properties.  Calcium  is  a  silvery  white  metal,  but  acquires  a  slightly 
yellowish  tinge  upon  exposure  to  air,  owing  to  its  union  with  nitrogen. 
It  has  a  density  of  1.55  and  melts  at  780°.  It  combines  readily  with 
most  of  the  nonmetals,  often  with  evolution  of  light.  For  example,  it 
combines  with  the  elements  of  the  sulfur  family,  the  chlorine  family, 
and  the  nitrogen  family  (bismuth  excepted).  When  heated  to  ignition 
in  oxygen,  it  burns  with  dazzling  brilliancy.  When  burned  in  air, 
both  the  oxide  and  the  nitride  of  calcium  are  obtained.  It  reacts  with 
water  and  with  dilute  acids,  as  represented  in  the  following  equations  : 


Calcium  promises  to  become  a  useful  metal,  although  its  commercial 
applications  are  as  yet  rather  limited.  It  is  a  powerful  reducing  agent 
and  would  undoubtedly  find  application  in  the  reduction  of  some  of 
the  metallic  oxides  if  it  could  be  produced  at  a  sufficiently  low  cost. 
A  limited  amount  of  it  is  used  in  the  preparation  of  the  hydride. 
Because  of  its  affinity  for  nitrogen  it  has  also  been  used  to  remove 
nitrogen  in  the  preparation  of  argon. 

Compounds  of  calcium.  The  preparation  of  the  compounds  of  cal- 
cium does  not  in  general  present  as  great  a  problem  as  does  the 
preparation  of  those  of  sodium.  This  is  due  to  the  fact  that  the  form 
in  which  calcium  occurs  most  abundantly  in  nature  is  the  carbonate, 
a  compound  which  is  readily  changed  into  other  compounds  by  the 
action  of  the  appropriate  acids.  Moreover,  the  carbonate  can  be  de- 
composed by  heat  without  difficulty,  thus  furnishing  an  inexpensive 
method  for  the  preparation  of  the  oxide  (lime). 

Since  the  compounds  found  in  nature  are  in  general  more  or  less  impure,  the 
chemist  must  find  some  method  for  removing  the  impurities  if  he  wishes  to  pre- 
pare chemically  pure  compounds  from  minerals.  A  number  of  different  methods 
of  procedure  are  used  in  such  cases,  the  choice  depending  upon  the  properties 
of  the  natural  compounds  and  the  character  of  the  impurities.  In  some  cases 
it  is  best  to  separate  the  pure  metal  by  electrolytic  methods  ;  in  other  cases  it 
is  more  feasible  to  decompose  the  mineral  and  separate  the  impurities  by  pre- 
cipitation. For  example,  suppose  we  wish  to  prepare  pure  calcium  carbonate 
from  limestone.  The  chief  impurities  present  in  the  limestone  are  compounds 


416  GENEBAL  CHEMISTRY 

of  iron,  aluminium,  and  magnesium.  The  limestone  is  dissolved  in  hydrochloric 
acid,  and  any  insoluble  matter  which  may  be  present  is  filtered  off.  The  filtrate 
contains  the  calcium,  together  with  the  other  metals  in  the  limestone  in  the 
form  of  chlorides.  By  adding  a  solution  of  calcium  hydroxide,  the  iron,  alu- 
minium, and  magnesium  are  all  precipitated  in  the  form  of  hydroxides  and  are 
removed  by  filtration.  The  filtrate  then  contains  calcium  chloride  and  calcium 
hydroxide.  If  a  solution  of  ammonium  carbonate  is  now  added  to  this  solution, 
pure  calcium  carbonate  is  precipitated. 

Calcium  oxide  (quicklime)  (CaO).    Calcium  oxide  may  be  obtained 
v   by  burning  the  metal  in  air  or  by  heating  the  nitrate  or  carbonate.   It 
is  obtained  commercially  by  heating  limestone  in  large  furnaces  called 
limekilns :  CaCO        >  CaO  -f  CO 

The  reaction  is  reversible,  as  is  indicated  in  the  equation.  If  the  decomposi- 
tion is  carried  out  in  a  closed  vessel,  equilibrium  between  the  opposing  reactions  is 
reached  at  any  definite  temperature  when  the  carbon  dioxide  evolved  at  that 
temperature  exerts  a  certain  pressure.  The  higher  the  temperature  the  greater 
the  pressure  at  equilibrium,  as  shown  in  the  following  table,  in  which  the  pressure 
is  expressed  in  millimeters  of  mercury : 

Temperature,    547°     610°     625°     740°     745°     810°     812°       865° 
Pressure,  27        46        56      255      289      678      753      1333 

At  any  given  temperature,  therefore,  say  740°,  calcium  carbonate  will  be  de- 
composed if  the  pressure  exerted  by  the  carbon  dioxide  is  less  than  255  mm.,  or 
it  will  be  formed  by  the  combination  of  calcium  oxide  and  carbon  dioxide  if  the 
pressure  is  greater  than  255  mm.,  the  reaction  proceeding  in  either  case  until  the 
pressure  equals  255  mm.,  when  equilibrium  results.  It  is  evident  that  the  reaction 
will  never  reach  completion  if  carried  out  in  a  closed  vessel.  If,  on  the  other  hand, 
the  carbon  dioxide  is  conducted  away  as  fast  as  formed,  the  decomposition  will  con- 
tinue until  complete.  This  is  the  method  adopted  in  the  production  of  lime,  the 
limestone  being  heated  in  a  current  of  air,  which  carries  away  the  carbon  dioxide. 

Pure  calcium  oxide  is  a  white,  amorphous  substance,  the  density 
of  different  specimens  varying  from  3  to  3.3.  When  heated  intensely, 
as  in  the  oxyhydrogen  blowpipe,  it  gives  a  brilliant  light  called  the 
limelight.  Although  it  is  a  substance  very  difficult  to  fuse,  it  not 
only  melts  but  boils  vigorously  at  the  temperature  of  the  electric 
furnace  (about  3500°).  Water  acts  upon  lime,  with  the  evolution 
of  considerable  heat,  the  process  being  called  slaking: 

CaO  +  H20  =  Ca(OH)2 

When  exposed  to  air,  calcium  oxide  is  gradually  converted  into  the 
hydroxide  and  carbonate  and  will  no  longer  slake  upon  addition  of 
water.  It  is  then  said  to  be  air-slaked.  Lime  is  produced  in  enormous 
quantities  and  is  used  in  making  calcium  hydroxide. 


THE  ALKALINE  EARTH  METALS 


417 


Commercial  production  of  lime.  The  older  form  of  kiln,  still  in  use,  consists  of 
a  large  stack,  or  chimney,  in  which  the  limestone  is  loosely  packed.  A  fire  is 
built  at  the  base  of  the  stack,  and  when  the  decomposition  of  the  limestone 
is  complete,  the  fire  is  allowed  to  die  out  and  the  lime  is  removed.  A  longitudinal 
section  of  the  newer  form  of  kiln  is  shown  in  Fig.  137.  The  kiln  is  about  50  ft.  in 
height.  A  number  of  fire  boxes,  or  furnaces,  A,  A,  are  built  around  the  lower  part, 
all  leading  into  the  central  stack.  The  kiln 
is  filled  with  limestone  through  a  swinging 
door  B.  The  hot  products  of  combustion 
are  drawn  up  through  the  kiln,'  and  the 
limestone  is  gradually  decomposed  by  the 
heat.  The  bottom  of  the  furnace  is  so  con- 
structed that  a  current  of  air  is  drawn  in 
at  C,  and  this  serves  the  double  purpose 
of  cooling  the  hot  lime  at  the  base  of  the 
furnace  and  furnishing  heated  oxygen  for 
the  combustion.  The  lime  is  dropped  into 
cars  run  under  the  furnace.  The  advantage 
of  this  kind  of  kiln  over  the  older  form  is 
that  the  process  is  continuous,  limestone 
being  charged  in  at  the  top  as  fast  as  the 
lime  is  removed  at  the  bottom. 

Calcium  hydroxide  (slaked  lime) 
(Ca(OH)2).  This  compound  is  pre- 
pared by  adding  water  to  calcium 
oxide.  Considerable  heat  is  devel- 
oped in  the  reaction,  as  is  repre- 
sented in  the  following  equation: 

CaO  +  HO  =  Ca(OH)2  + 15,540  cal. 


FIG.  137 


The  reaction  is  reversible,  the  decom- 
position  of  the    hydroxide    into  its 

components  taking  place  rapidly  when  heated  in  an  open  vessel  to  a 
temperature  above  450°.  Pure  calcium  hydroxide  is  a  white  powder 
having  a  density  of  about  2.  It  is  but  slightly  soluble  in  water,  its 
solubility  diminishing  with  rise  in  temperature.  Thus,  at  10°,  1  1.  of 
water  dissolves  1.76 g.  of  the  hydroxide;  at  50°,  1.28 g.;  and  at  100°, 
only  0.77  g.  Its  solution  in  water  is  termed  limewater  and  is  often 
used  in  medicine  because  of  its  basic  properties.  Calcium  hydroxide  is 
a  moderately  strong  base  (p.  155)  and,  owing  to  its  cheapness,  is  much 
used  in  the  industries  whenever  an  alkali  is  desired.  A  number  of 
its  uses  have  already  been  mentioned.  It  is  used  in  the  preparation 
of  ammonia,  bleaching  powder,  and  the  hydroxides  of  sodium  and 


418  GENERAL  CHEMISTRY 

potassium.  It  is  also  used  in  the  purification  of  coal  gas  (p.  323),  in 
removing  the  hair  from  hides,  in  the  manufacture  of  leather,  and  in 
making  mortar  and  plaster. 

Mortar  is  a  mixture  of  calcium  hydroxide  and  sand.  When  it  is  exposed  to 
the  air  or  spread  upon  porous  materials,  moisture  is  removed  from  it,  partly  by 
absorption  in  the  porous  materials  and  partly  by  evaporation,  and  the  mortar 
becomes  firm,  or  sets.  At  the  same  time  carbon  dioxide  is  slowly  absorbed  from 
the  air,  and  hard  calcium  carbonate  is  formed : 

Ca(OH)2  +  C02  =  CaC08  +  H2O 

By  this  combined  action  the  mortar  becomes  very  hard  and  adheres  firmly  to  the 
surface  upon  which  it  is  spread.  The  sand  serves  to  give  body  to  the  mortar 
and  makes  it  porous.  It  also  prevents  too  much  shrinkage. 

Sulfides  of  calcium.  The  normal  sulfide  (CaS)  is  prepared  by  re- 
ducing calcium  sulfate  with  carbon.  It  is  obtained  as  a  by-product 
in  the  Leblanc  process  for  the  manufacture  of  sodium  carbonate. 
Although  insoluble  in  water,  it  gradually  undergoes  hydrolysis,  form- 
ing the  acid  sulfide  (Ca(HS)2),  which  is  soluble.  A  mixture  of  the 
tetrasulfide  (CaS4)  and  pentasulfide  (CaS5)  can  be  obtained  by  heat- 
ing calcium  hydroxide  and  sulfur  in  the  presence  of  water  (p.  211). 

The  normal  sulfide  as  prepared  commercially  is  sometimes  used  as 
a  pigment  for  luminous  paint,  since  after  exposure  to  a  bright  light  it 
will  glow  in  the  dark.  It  is  interesting  to  note  that  the  pure  sulfide 
does  not  possess  this  property,  but  acquires  it  in  the  presence  of  small 
percentages  of  the  sulfides  of  some  other  metals,  especially  those  of 
manganese,  bismuth,  and  vanadium. 

Calcium  fluoride  (CaF2).  This  salt  occurs  in  large  quantities  in 
nature  as  flu  or  spar.  The  mineral  crystallizes  in  the  form  of  cubes  or 
octahedra,  and  large  crystals  are  often  found  that  are  beautifully 
tinted,  generally  a  shade  of  green  or  blue.  When  heated  gently  they 
become  fluorescent.  Calcium  fluoride  also  occurs  in  the  ashes  of  some 
plants.  When  pure  it  is  a  white  solid  and  is  nearly  insoluble  in  water. 
It  melts  at  1330°.  The  mineral  fluor  spar  is  mined  in  large  quantities, 
especially  in  southern  Illinois,  and  increasing  amounts  of  it,  are  being 
used  as  a  flux  in  various  metallurgical  operations.  It  is  also  used  in 
the  manufacture  of  opaque  glass  and  white  enamel  and  in  the  prepa- 
ration of  other  fluorine  compounds,  especially  hydrofluoric  acid. 

Calcium  chloride  (CaCl2).  This  salt  is  present  in  sea  water  to  a 
limited  extent.  Small  quantities  also  occur  as  a  constituent  of  some 
minerals,  such  as  tachhydrite  (2  MgCl2  •  CaCl2  •  12  H2O),  found  in  the 


THE  ALKALINE  EAETH  METALS         419 

Stassfurt  deposits.  It  is  formed  in  large  quantities  as  a  by-product  in 
the  manufacture  of  sodium  carbonate  by  the  Solvay  process,  as  well  as 
in  the  preparation  of  potassium  chlorate  from  calcium  chlorate.  The 
mother  liquor  from  salt  works  also  contains  a  considerable  percent- 
age of  the  compound.  The  supply  of  the  compound  obtained  from 
these  sources  is  at  present  much  greater  than  the  demand.  Pure  cal- 
cium chloride  may  be  prepared  by  dissolving  calcium  carbonate  in 
hydrochloric  acid  and  evaporating  to  crystallization.  It  separates  from 
saturated  solutions  at  ordinary  temperatures  in  the  form  of  hexagonal 
crystals  which  have  -the  formula  CaCl2*6H2O.  The  mono-,  di-,  and 
tetra-hydrates  have  also  been  prepared.  When  the  hydrates  are  heated 
above  260°,  the  anhydrous  salt  is  obtained  as  a  white,  porous  mass.  In 
this  process  some  calcium  oxide  is  formed  by  the  action  of  water  upon 
thechloride:  " 


As  a  rule,  therefore,  a  solution  of  the  anhydrous  salt  reacts  basic,  owing 
to  the  presence  of  calcium  hydroxide. 

The  anhydrous  salt  readily  absorbs  moisture  and  is  largely  used  as 
a  drying  agent.  It  melts  at  780°.  It  dissolves  in  water  with  the  lib- 
eration of  considerable  heat,  while  the  hexahydrate  absorbs  heat  under 
the  same  conditions.  It  is  very  soluble,  100  g.  of  water  dissolving 
59.5  g.  at  0°  ;  74.5  g.  at  20°  ;  and  159  g.  at  100°.  Owing  to  its  great 
solubility,  the  boiling  point  of  a  concentrated  solution  of  calcium 
chloride  is  much  higher  than  that  of  water  ;  thus,  a  solution  contain- 
ing 101  g.  of  the  chloride  in  100  g.  of  water  boils  at  130°.  The 
hexahydrate  and  ice  form  a  freezing  mixture  by  means  of  which  a 
temperature  as  low  as  —  50°  can  be  reached.  A  solution  of  calcium 
chloride  is  largely  used  as  a  brine  in  the  manufacture  of  artificial  ice 
(p.  174).  Because  of  its  deliquescent  character  it  has  been  used  to  lay 
the  dust  on  roads.  Mines  have  also  been  sprinkled  with  its  solution 
in  the  hope  of  preventing  dust  explosions  (p.  331). 

Calcium  carbonate  (CaC03).  Enormous  quantities  of  calcium  car- 
bonate occur  in  nature.  Limestone  is  the  most  abundant  form  and  is 
a  grayish  rock  usually  found  in  hard,  stratified  masses.  Whole  moun- 
tain ranges  are  sometimes  made  up  of  this  material.  It  is  never  pure 
calcium  carbonate,  always  containing  variable  percentages  of  foreign 
matter,  usually  magnesium  carbonate,  clay,  silica,  compounds  of  iron 
and  aluminium,  and  frequently  fossil  remains.  Marl  is  a  mixture  of 
limestone  and  clay.  Pearls,  coral,  and  various  kinds  of  shells,  such  as 


420  GENERAL  CHEMISTRY 

eggshells  and  oyster  shells,  and  natural  chalk  (not  the  blackboard 
crayon)  are  largely  calcium  carbonate. 

Calcium  carbonate  is  a  dimorphous  substance.  The  more  common 
form  of  crystal  belongs  to  the  hexagonal  system.  Calcite  is  a  pure, 
crystalline  form  of  this  character.  Very  beautiful  transparent  crystals 
of  calcite  are  found  in  Iceland ;  hence  the  name  Iceland  spar,  often 
applied  to  this  variety.  Mexican  onyx  is  a  massive  variety  streaked 
or  banded  with  colors  occasioned  by  impurities.  Marble  is  made  up 
of  minute  snow-white  calcite  crystals  and  was  probably  formed  by 
the  crystallization  of  the  melted  rock  under  great  pressure.  Stalactites 
and  stalagmites  are  icicle-like  forms  sometimes  found  in  caves. 

Calcium  carbonate  is  also  found  in  nature  in  the  form  of  crystals 
belonging  to  the  rhombic  system.  The  mineral  aragonite  is  an  example 
of  this  form.  When  heated  it  is  transformed  into  a  mass  of  calcite 
crystals.  This  transformation  is  probably  taking  place  at  ordinary 
temperatures  but  with  a  very  low  speed.  It  is  easy  to  understand, 
therefore,  why  the  aragonite  form  is  a  rather  uncommon  variety. 

Preparation  and  uses  of  calcium  carbonate.  In  the  laboratory  pure 
calcium  carbonate  can  be  prepared  by  treating  a  solution  of  a  calcium 
salt,  such  as  the  chloride,  with  a  solution  of  a  carbonate : 

CaCl2  +  Na2CO8  =  CaCO3  +  2  NaCl 

The  insoluble  carbonate  precipitates  and  may  be  separated  by  filtration. 
When  the  precipitation  is  effected  at  ordinary  temperatures,  the  carbo- 
nate is  obtained  as  an  amorphous  powder  which  soon  changes  over  into 
calcite  crystals.  When  precipitated  from  hot  dilute  solutions,  aragonite 
crystals  are  obtained,  but  these  slowly  change  into  the  calcite  form. 

Pure  calcium  carbonate  is  very  slightly  soluble  in  water,  1 1.  of  water 
at  15°  dissolving  but  13  mg.,  or  if  it  is  in  the  form  of  aragonite,  15  mg. 
The  action  of  acids  upon  this  substance,  as  well  as  the  effect  of  heat 
upon  it,  has  already  been  discussed. 

The  natural  varieties  of  calcium  carbonate  find  many  uses,  such  as 
in  the  preparation  of  lime,  cement,  and  carbon  dioxide,  in  metallurgi- 
cal operations  (especially  in  blast  furnaces),  in  the  manufacture  of 
soda  and  glass,  and  for  building  stone  and  ballast  for  roads. 

Calcium  acid  carbonate  (calcium  bicarbonate)  Ca(HC03)2.  While 
calcium  carbonate  is  almost  insoluble  in  pure  water,  it  readily  dis- 
solves in  water  containing  carbon  dioxide.  This  is  undoubtedly  due 
to  the  formation  of  the  soluble  but  unstable  calcium  acid  carbonate. 


THE  ALKALINE  EAKTH  METALS         421 

When  solutions  containing  the  acid  carbonate  are  heated,  the  normal 
carbonate  is  precipitated  (p.  286) : 

Ca(HC08)2  =  CaC08  +  H2O  +  CO2 

Natural  waters  always  contain  more  or  less  carbon  dioxide  in  solution.  In 
the  case  of  certain  underground  waters  the  amount  of  carbon  dioxide  is  compar- 
atively large,  being  held  in  solution  by  pressure.  Such  waters  have  a  marked 
solvent  action  upon  limestone,  dissolving  both  the  calcium  carbonate  and  the 
magnesium  carbonate.  In  certain  localities  this  solvent  action,  continued  through 
geological  ages,  has  resulted  in  the  formation  of  large  caves  in  limestone  rock, 
such  as  the  Mammoth  Cave  in  Kentucky. 

Calcium  sulfate  (CaSOJ.  This  compound  occurs  in  nature  in  the 
anhydrous  form  in  the  mineral  anhydrite.  More  commonly,  however, 
it  is  found  as  gypsum,  which  is  a  dihydrate  (CaSO4  •  2  H2O).  Several 
other  forms  of  the  dihydrate  are  also  known,  such  as  selenite,  alabaster, 
and  satin  spar.  When  the  dihydrate  is  heated,  there  is  formed  a  hydrate 
of  the  formula  2  CaSO4  •  H2O.  Since  this  compound  contains  half  as 
many  molecules  of  water  of  crystallization  as  of  calcium  sulfate,  it 
is  called  a  hemihydrate. 

Of  the  different  forms  of  calcium  sulfate,  gypsum  is  by  far  the 
most  important  commercially.  It  is  quarried  in  large  quantities,  espe- 
cially in  New  York,  Michigan,  and  Oklahoma.  It  crystallizes  in  six- 
sided  prisms  belonging  to  the  monoclinic  system.  It  is  but  slightly 
soluble  in  water,  its  solubility  increasing  slowly  from  0°  to  about  35° 
and  then  decreasing.  It  is  used  in  making  plaster  of  Paris,  as  a  filler  in 
making  paper,  as  a  paint  pigment,  and  as  a  constituent  of  fertilizers. 
Its  value  as  a  fertilizer  seems  to  be  due  to  the  fact  that  it  reacts  with 
the  ammonium  carbonate  present  in  the  soil,  forming  ammonium  sul- 
fate, which  is  much  less  volatile  than  the  carbonate  and  is  therefore 
retained  in  the  soil  until  taken  up  by  the  growing  plants. 

Plaster  of  Paris.  This  is  a  fine  white  powder  obtained  by  heating 
gypsum,  and  consists  essentially  of  the  hemihydrate  2  CaSO4  •  H2O. 
When  water  is  added,  this  powder  forms  a  plastic  mass,  which  quickly 
hardens,  or  sets,  and  regains  its  crystalline  structure.  These  prop- 
erties make  it  valuable  as  a  material  for  molding  casts  and  stucco 
work,  for  cementing  glass  to  metals,  and  for  a  finishing  coat  on 
plastered  walls.  In  the  manufacture  of  plaster  of  Paris  the  tempera- 
ture must  not  be  allowed  to  rise  much  above  125° ;  otherwise  the 
anhydrous  salt  is  formed,  and  this  combines  with  water  so  slowly  as  to 
render  it  worthless  for  the  purposes  for  which  plaster  of  Paris  is  used. 


422  GENERAL  CHEMISTEY 

Hard  water.  Waters  containing  compounds  of  calcium  and  mag- 
nesium in  solution  are  called  hard  waters.  The  hardness  of  water 
may  be  of  two  kinds :  (1)  temporary  hardness,  and  (2)  permanent 
hardness. 

1.  Temporary  hardness.  We  have  seen  that  when  water  charged  with 
carbon  dioxide  comes  in  contact  with  limestone,  a  certain  amount  of 
the  latter  dissolves,  owing  to  the  formation  of  the  soluble  acid  car- 
bonate of  calcium.    The  hardness  of  such  waters  is  said  to  be  tempo- 
rary, since  it  may  be  removed  by  boiling.    The  heat  changes  the  acid 
carbonate  into  the  insoluble  normal  carbonate  which  then  precipitates, 
rendering  the  water  soft : 

Ca(HC03)2  =  CaC03  +  H2O  +  CO2 

Such  waters  may  also  be  softened  by  the  addition  of  sufficient  lime 
or  calcium  hydroxide  to  convert  the  acid  carbonate  of  calcium  into 
the  normal  carbonate : 

Ca(HC03)2  +  Ca(OH)2  =  2  CaCO3  +  2  H2O 

2.  Permanent  hardness.  The  hardness  of  water  may  also  be  due  to 
the  presence  of  calcium  and  magnesium  sulfates  or  chlorides.   Boiling 
the  water  does  not  affect  these  salts;  hence  such  waters  are  said  to 
have  permanent  hardness.    They  may  be  softened,  however,  by  the 
addition   of  sodium  carbonate,   which  precipitates  the  calcium  and 
magnesium  as  insoluble  carbonates : 

CaS04  +  Na2C03  -  CaCO3  +  Na2SO4 
This  process  is  sometimes  called  "  breaking  "  the  water. 

Commercial  methods  for  softening  water.  The  average  water  of  a  city  supply 
contains  not  only  the  acid  carbonates  of  calcium  and  magnesium  but  also  the 
sulfates  and  chlorides  of  these  metals,  together  with  other  salts  in  smaller  quan- 
tities. Such  waters  are  softened  on  a  commercial  scale  by  the  addition  of  the 
proper  quantities  of  calcium  hydroxide  and  sodium  carbonate.  The  calcium 
hydroxide  precipitates  the  acid  carbonates,  while  the  sodium  carbonate  precipi- 
tates the  other  soluble  salts  of  calcium  and  magnesium.  The  amounts  of  calcium 
hydroxide  and  sodium  carbonate  required  to  soften  any  given  water  are  calcu- 
lated from  a  chemical  analysis  of  the  water.  It  will  be  noticed  from  the  equa- 
tions that  the  water  softened  in  this  way  contains  sodium  sulfate  and  chloride, 
but  the  presence  of  these  salts  is  not  objectionable. 

Sulfites  of  calcium.  The  normal  sulfite  CaSOg,  being  but  slightly 
soluble,  is  formed  as  a  white  precipitate  when  a  solution  of  sodium  sul- 
fite is  added  to  a  solution  of  a  calcium  salt.  When  it  is  suspended  in 


THE  ALKALINE  EARTH  METALS 


423 


water  and  sulfur  dioxide  is  passed  into  the  mixture,  it  dissolves,  owing 
to  the  formation  of  the  soluble  bisulfite  Ca(HSO3)0.  This  solution 
is  used  as  a  preservative,  and  in  much  larger  quantities  in  the  manu- 
facture of  paper  (p.  216).  For  this  purpose  it  is  prepared  directly  by 
passing  sulfur  dioxide  into  a  solution  of  calcium  hydroxide : 

Ca(OH)2  +  2  SO2  =  Ca(HSO3)2 

Calcium  carbide  (CaC2).  This  compound  is  now  prepared  on  a  large 
scale  for  use  in  generating  acetylene  (p.  295)  and  in  making  fertilizer. 
It  was  first  obtained  in  impure  form  by  E.  Davy  in  1836,  and  by 
Wohler  in  1863.  Later,  Borchers  and  also  Moissan  obtained  it  in  pure 
crystalline  form.  In  1893  Wilson  devised  a  method  for  preparing  it 
on  a  large  scale,  and  since  that  time  it  has  been  a  commercial  product. 
It  is  made  by  heating  a  mixture  of  coke  and  lime  in  an  electric  furnace : 

CaO  +  3  C  =  CaC2  +  CO 

A  large  amount  of  heat  (12,000  cal.)  is  absorbed  in  the  reaction,  and 
the  carbide  is  strongly  endothermic. 

The  pure  carbide  is  a  colorless,  transparent,  crystalline  substance 
and  is  practically  insoluble  in  all  known  solvents.  At  high  tempera- 
tures it  is  a  powerful  reducing  agent.  Its  commercial  importance  lies 
in  the  fact  that  it  reacts  with  water 
to  form  acetylene  and  with  nitro- 
gen to  form  cyanamide  (CaN  C). 


CaC 


H2O  =  CaO 


C2H2 


The  commercial  article  is  a  dull 
gray,  porous  substance  which  con- 
tains many  impurities.  The  acetylene 
prepared  from  this  substance  has  a 
disagreeable  odor  due  to  phosphine 
and  other  impurities. 

Commercial  production  of  calcium  car- 
bide. A  number  of  different  forms  of  fur- 
naces are  used.  The  general  principles 
involved,  however,  may  be  illustrated  by 
the  diagram  shown  in  Fig.  138,  which  rep- 

resents a  simple  type  of  these  furnaces.  The  base  of  the  furnace  is  provided  with 
a  large  block  of  carbon  A,  which  serves  as  one  of  the  electrodes.  The  other  elec- 
trodes B,  B,  several  in  number,  are  arranged  horizontally  at  some  distance  above 


FIG.  138 


424  GENERAL  CHEMISTRY 

this.  A  mixture  of  coke  and  lime  is  fed  into  the  furnace  through  the  trap  top  C. 
An  alternating  current  is  used,  and  this  is  regulated  so  as  to  give  a  temperature 
of  about  2000°.  At  this  temperature  the  carbide  is  formed  and  settles  to  the 
bottom  of  the  furnace  in  liquid  state.  This  is  drawn  oft'  through  the  tap  hole  D. 
The  carbon  monoxide  generated  in  the  reaction  escapes  through  the  pipes  E,  E, 
and  is  led  back  into  the  furnace.  The  pipes  F,  F  supply  air,  so  that  the  monoxide 
burns  as  it  reenters  the  furnace  and  assists  in  heating  the  charge.  The  carbon 
dioxide  so  formed,  together  with  the  nitrogen  entering  in  the  air,  escapes  at  G. 

Calcium  cyanamide  (CaN2C).  This  compound  is  formed  by  passing 
nitrogen  over  calcium  carbide.  The  reaction  takes  place  best  at  about 
1200°.  The  compound  is  a  derivative  of  cyanamide  (H2NNC),  the 
two  hydrogen  atoms  in  each  molecule  being  replaced  by  a  divalent 
calcium  atom.  When  heated  with  carbon  it  is  changed  into  calcium 
cyanide.  Its  chief  importance,  however,  lies  in  the  fact  that  it  is  a 
nitrogenous  fertilizer,  all  of  its  nitrogen  being  available  for  absorption 
by  growing  plants. 

In  the  commercial  preparation  of  calcium  cyanamide  the  nitrogen  is  obtained 
by  passing  air  over  sodium  hydroxide  to  remove  the  moisture  and  carbon  dioxide, 
and  then  over  red-hot  copper  to  remove  the  oxygen.  The  nitrogen  so  obtained 
is  passed  over  the  carbide  packed  in  suitable  tubes.  Through  the  central  portion 
of  the  carbide  is  a  carbon  rod.  The  heat  necessary  for  the  reaction  is  generated 
by  passing  an  electric  current  through  this  rod.  Inasmuch  as  the  carbide  used  is 
very  impure,  the  product  contains  only  about  60  per  cent  of  cyanamide,  the  other 
chief  ingredients  being  lime  and  carbon.  This  product  is  known  as  lime  nitrogen. 
It  is  ground  and  treated  with  water,  which  slakes  the  lime,  and  in  this  form  is 
sold  as  a  fertilizer  under  the  name  cyanamide. 

The  utilization  of  atmospheric  nitrogen.  It  has  been  pointed  out  that,  with  few 
exceptions,  organisms  have  not  the  power  of  directly  assimilating  free  nitrogen 
(p.  107).  Repeated  attempts  have  therefore  been  made  to  utilize  the  inexhaust- 
ible supplies  of  free  nitrogen  in  the  atmosphere  by  converting  the  nitrogen  into 
compounds  which  contain  the  element  in  a  form  available  for  plant  food.  The 
following  methods  may  be  used  for  effecting  this  change :  (1)  the  nitrogen  may 
be  converted  into  calcium  cyanamide  as  described  above  ;  (2)  the  nitrogen  may  be 
converted  into  nitric  acid  and  then  into  nitrates  (p.  177) ;  (3)  ammonia  may 
be  formed  by  heating  a  mixture  of  nitrogen  and  hydrogen  under  high  pressure 
(200  atmospheres)  and  in  contact  with  a  suitable  catalytic  agent,  such  as  finely 
divided  iron ;  (4)  nitrides  of  certain  metals,  such  as  aluminium,  may  be  formed 
by  the  direct  union  of  the  two  elements,  and  from  these  nitrides  ammonia  may 
be  generated  through  the  action  of  steam. 

The  first  and  second  of  these  methods  are  now  used  commercially,  and  it  is 
claimed  that  the  third  has  been  developed  to  an  extent  that  likewise  insures  its 
economic  success.  It  seems  certain  that  the  compounds  so  formed,  or  similar  ones 
derived  from  atmospheric  nitrogen,  will  eventually  replace  the  sodium  nitrate 
and  the  ammonium  salts  which  are  now  the  chief  nitrogenous  products  used 
in  the  manufacture  of  fertilizers. 


THE  ALKALINE  EARTH  METALS         425 

Phosphates  of  calcium.  With  phosphoric  acid,  calcium  forms  three 
salts,  the  names  and  formulas  of  which  are  as  follows: 

Normal  (or  tertiary)  calcium  phosphate    .     .     Ca3(PO4)2 
Primary  calcium  phosphate    .     .  '  .     .     .     .     Ca(H2PO4)2 
Secondary  calcium  phosphate      .     ,     .     .     .     CaHPO4 

The  normal  phosphate,  usually  called  simply  calcium  phosphate,  is 
found  in  quantities  in  nature,  largely  in  the  form  of  phosphorite.  It 
is  the  chief  mineral  constituent  of  bones,  the  ash  of  which  contains 
about  80  per  cent  of  this  compound.  It  can  be  obtained  by  adding 
ammonium  hydroxide  to  a  solution  of  a  calcium  salt  until  strongly 
alkaline  and  then  precipitating  with  disodium  phosphate  : 

3  CaCl2  +  2  Na2HP04  +  2  NH-OH 

=  Ca3(PO4)2  +  4  NaCl  +  2  NH4C1  +  2  H2O 

It  is  nearly  insoluble  in  water  but  easily  dissolves  in  acids,  even  in  very 
weak  ones  like  acetic.  The  importance  of  the  phosphates  in  connection 
with  the  subject  of  fertilizers  has  already  been  discussed  (p.  363). 

Primary  calcium  phosphate  is  deposited  in  the  form  of  white  crystals  when 
a  solution  of  the  normal  phosphate  in  phosphoric  acid  is  evaporated  : 

Ca3(P04)2  4-  4  H3P04  =  3  Ca(H2PO4)2 

When  a  solution  of  disodium  phosphate  is  added  to  a  solution  of  a  calcium 
salt,  secondary  calcium  phosphate  is  obtained  : 

CaCl2  +  Na2HPO4  =  CaHPO4  +  2  XaCl 


Both  of  the  acid  salts  are  also  formed,  along  with  calcium  sulf  ate,  by  the  action 
of  sulf  uric  acid  upon  normal  calcium  phosphate  (p.  364). 

Silicates  of  calcium.  A  number  of  these  are  known.  The  metasili- 
cate  CaSiO3  occurs  pure  in  nature  in  the  form  of  the  mineral  wollaston- 
ite.  Combined  with  the  silicates  of  other  metals,  calcium  silicates  are 
widely  distributed.  They  can  be  prepared  by  fusing  lime  and  silica 
(sand)  together.  They  derive  their  chief  interest  from  the  fact  that 
they  are  important  constituents  of  cement  and  glass. 

Calcium  oxalate  (CaC204).  This  compound  owes  its  chief  interest 
to  the  fact  that  it  is  one  of  the  most  insoluble  of  the  compounds 
of  calcium.  When  a  soluble  oxalate,  such  as  ammonium  oxalate 
((NH4)0C2O4),  is  added  to  a  neutral  solution  of  any  calcium  compound, 
calcium  oxalate  precipitates  as  a  fine  white  powder.  A  solution  of 
ammonium  oxalate  is  therefore  used  as  a  reagent  for  the  detection 
and  estimation  of  calcium. 


426  GENERAL  CHEMISTRY 

STRONTIUM 

Occurrence  and  preparation.  Although  somewhat  widely  distributed, 
strontium  is  the  least  abundant  of  the  alkaline  earth  metals.  Its  chief 
minerals  are  celestite  (SrSO4)  and  strontianite  (SrCO3).  The  former 
is  the  more  abundant  and  is  found  especially  on  some  of  the  islands 
in  Lake  Erie.  At*  Put-in-Bay  beautiful  large  crystals  of  the  mineral 
are  found,  lining  the  walls  of  Strontia  Cave. 

The  metal  was  first  isolated  by  Davy  (1807),  although  in  an  impure 
state.  Its  preparation  is  much  more  difficult  than  that  of  calcium, 
although  the  general  methods  involved  are  the  same.  It  closely 
resembles  calcium  in  its  general  properties. 

Compounds  of  strontium.  Celestite  serves  as  the  source  material  for 
the  preparation  of  the  other  compounds.  Since  this  is  but  slightly 
soluble  and  is  not  acted  upon  by  acids  to  any  extent,  it  is  first 
converted  into  the  sulfide  by  reduction  with  carbon,  or  into  the 
carbonate  by  fusion  with  sodium  carbonate.  These  two  compounds, 
being  salts  of  volatile  acids,  are  readily  changed  into  other  salts  by 
the  action  of  the  appropriate  acids.  The  compounds  of  strontium  are 
very  similar  to  those  of  calcium  and  for  this  reason  will  be  discussed 
very  briefly. 

Strontium  oxide  (SrO)  ;  strontium  hydroxide  (Sr(OH)2).  The  oxide 
is  obtained  by  heating  the  carbonate.  It  combines  with  water  to  form 
the  hydroxide,  which  is  a  moderately  strong  base.  The  latter  is  fairly 
soluble  in  hot  water  and  crystallizes  from  the  solution  in  the  form  of 
the  octahydrate  Sr(OH)2  •  8  H2O.  The  hydroxide  forms  with  sucrose 
an  insoluble  compound  which  can  easily  .be  decomposed  into  its  orig- 
inal components.  It  has  therefore  been  used  in  the  refining  of  sugar, 
to  extract  the  sugar  from  uncrystallizable  sirups. 

Strontium  nitrate  (Sr(N03)2).  This  salt  separates  from  a  hot,  aqueous 
solution  in  the  anhydrous  form,  while  from  a  cold  solution  it  separates 
as  a  tetrahydrate  Sr(NOg)2  •  4  H2O.  When  ignited  with  combustible 
materials  it  imparts  a  brilliant  crimson  color  to  the  flame  and  is 
therefore  used  in  the  manufacture  of  red  lights. 

Other  compounds  of  strontium.  Among  the  other  compounds 
of  strontium  the  following  may  be  mentioned:  strontium  chloride 
(SrCl2*4H2O),  strontium  bromide  (SrBr2),  and  strontium  iodide  (SrI2), 
which  are  all  white  solids,  very  soluble  in  water;  strontium  sulfide 
(SrS),  which  is  even  more  phosphorescent  than  calcium  sulfide  under 


THE  ALKALINE  EAKTH  METALS         427 

the  same  conditions  ;  and  strontium  carbonate  (SrCO3),  which  occurs  in 
nature  as  strontianite  and,  being  insoluble,  can  be  prepared  from  the 
chloride  or  nitrate  by  precipitation  with  ammonium  carbonate. 

BARIUM 

Occurrence  and  preparation.  Like  strontium  and  calcium,  barium 
is  widely  distributed  as  a  constituent  of  igneous  rocks.  Its  most 
abundant  forms  are  barite  (or  barytes)  (BaSO4)  and  witherite  (BaCO^. 
The  former  is  the  more  abundant  and  is  mined  in  considerable  quan- 
tities, especially  in  Missouri,  for  use  as  a  paint  pigment.  The  state- 
ments made  in  reference  to  the  preparation  and  properties  of  strontium 
apply  equally  to  barium. 

Compounds  of  barium.  The  compounds  of  barium  are  very  similar 
to  those  of  calcium  and  strontium.  They  are  prepared  either  directly 
or  indirectly  from  barite.  This  is  first  converted  into  the  sulfide  or 
carbonate,  as  in  the  case  of  the  corresponding  strontium  sulfate,  and 
from  these  the  chloride  and  nitrate  are  prepared  by  the  action  of  the 
appropriate  acids.  The  soluble  compounds  of  barium  are  poisonous. 

Barium  oxide  (BaO)  ;  barium  peroxide  (Ba02).  Barium  oxide  is  pre- 
pared by  heating  the  nitrate  rather  than  the  carbonate,  since  the  latter 
is  decomposed  only  with  great  difficulty  : 

2  Ba(N03)2  -  2  BaO  +  4  NO2  +  O2 

Heated  to  a  low  red  heat  in  the  air,  barium  oxide  combines  with 
oxygen,  forming  the  peroxide  BaO2: 


This  reaction  is  reversible,  and  it  will  be  recalled  that  it  serves  as  a 
method  for  separating  oxygen  from  the  air  (p.  18). 

The  dissociation  of  barium  peroxide  into  barium  oxide  and  oxygen  is  in 
principle  exactly  like  the  dissociation  of  calcium  carbonate  into  calcium  oxide 
and  carbon  dioxide.  An  equilibrium  is  reached  between  the  opposing  reactions 
at  any  definite  temperature  when  the  oxygen  evolved  exerts  a  certain  pressure, 
as  given  in  the  following  table,  in  which  the  pressure  is  expressed  in  millimeters 
of  mercury  : 

Temperature,   525°    555°    650°    670°    720°    735°    750°    775°    790° 
Pressure,  20        25        65        80      210      260      340      510      670 

It  is  evident  that  at  any  definite  temperature  the  reaction  may  be  made  to  go  in 
either  direction  by  simply  varying  the  pressure.  In  the  Brin  process  for  the 
preparation  of  oxygen  it  was  found  more  economical  to  control  the  course  of  the 
reaction  by  variation  of  pressure  rather  than  by  variation  of  temperature. 


428  GENERAL  CHEMISTRY 

Barium  peroxide  is  a  white  solid,  insoluble  in  water.  It  combines 
with  water  to  form  the  hydrate  BaO2  •  10  H2O.  When  treated  with 
acids  it  yields  hydrogen  peroxide  (p.  70)  and  serves  for  the  commer- 
cial preparation  of  this  important  compound.  Strontium  and  calcium 
form  similar  peroxides,  but  they  are  not  so  readily  obtained. 

Barium  hydroxide  (Ba(OH)2).  Barium  oxide  resembles  the  oxides  of 
calcium  and  strontium  in  that  it  readily  combines  with  water  to  form 
the  corresponding  hydroxide.  The  hydroxide  forms  a  number  of  hy- 
drates, the  most  common  one  being  the  octahydrate  (Ba(OH)2  •  8  H2O), 
which  crystallizes  from  solutions  at  ordinary  temperatures.  It  is  much 
more  soluble  than  either  calcium  or  strontium  hydroxide,  and  its  solu- 
tion is  often  used  as  a  reagent  for  detecting  carbon  dioxide,  since  it 
forms  with  it  the  difficultly  soluble  barium  carbonate.  Its  aqueous 
solution  acts  as  a  strong  base. 

Barium  chloride  (BaCl2).  This  salt  is  prepared  by  the  action  of 
hydrochloric  acid  upon  barium  carbonate  or  sulfide.  It  is  quite  soluble 
in  water  and  crystallizes  from  saturated  solutions  in  the  form  of  white 
crystals  which  have  the  formula  BaCl2  •  2  H2O.  It  is  used  as  a  reagent 
for  the  detection  of  sulfuric  acid  or  a  soluble  sulfate,  forming  with 
them  the  insoluble  barium  sulfate  (p.  224). 

Barium  sulfate  (BaSOJ.  This  compound  has  been  known  for  a  long 
time  and  is  by  far  the  most  widely  used  of  the  compounds  of  barium. 
The  native  barium  sulfate  is  a  heavy  mineral ;  hence  the  name  barite, 
meaning  "  heavy,"  from  which  name  that  of  the  metal  itself  was  de- 
rived. Large  deposits  of  barite  occur  in  Missouri,  Nevada,  and  Cali- 
fornia. It  is  the  least  soluble  of  all  the  sulfates.  It  is  precipitated, 
even  in  the  presence  of  strong  acids,  when  a  solution  of  a  sulfate  (or 
sulfuric  acid)  is  added  to  a  solution  of  a  barium  salt.  The  native 
sulfate,  as  well  as  that  prepared  by  precipitation,  is  used  in  large 
quantities  as  a  pigment. 

Other  compounds  of  barium.  Barium  carbonate  (BaCO3)  is  sometimes 
mixed  with  clay  in  the  manufacture  of  terra-cotta  ware.  Barium 
nitrate  (Ba(NO8)2)  is  an  oxidizing  agent,  and  combustible  materials 
mixed  with  it  burn  with  a  green  flame.  It  is  therefore  used  in  the 
manufacture  of  green  lights.  Barium  sulfide  (BaS)  resembles  the  sul- 
fides  of  calcium  and  strontium  in  that  it  is  phosphorescent  and  hydro- 
lyzes  with  water,  forming  the  hydroxide  and  acid  sulfide.  Barium 
chr ornate  (BaCrO4)  is  a  yellow  solid,  insoluble  in  water. 


CHAPTER  XXXI 

THE  MAGNESIUM  FAMILY 


ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

BOILING 
POINT 

OXIDE 

Glucinum  (Gl)  .  .  . 
Magnesium  (Mg)  .  .' 
Zinc  (Zn)  

9.1 
24.32 
65.37 

1.93 
1.74 
7.10 

1430.° 
633.° 
419.4° 

1120° 
918° 

'     G1O 
MgO 
ZnO 

Cadmium  (Cd)  .  .  . 

112.4 

8.64 

321.° 

778° 

CdO 

The  family.  In  the  magnesium  family  are  included  the  five  elements, 
glucinum,  magnesium,  zinc,  cadmium,  and  mercury.  Among  the  first 
four  of  these  metals  there  is  a  close  family  resemblance,  such  as  has 
been  traced  between  the  members  of  the  two  preceding  families. 
Mercury  will  be  described  in  connection  with  copper,  since  in  some 
respects  it  is  more  like  that  metal  than  it  is  like  the  members  of  the 
magnesium  family. 

The  elements.  Like  the  metals  of  the  alkali  and  alkaline  earth 
families,  the  members  of  the  magnesium  family  exhibit  a  somewhat 
regular  gradation  in  properties.  Their  densities  (magnesium  excepted) 
are  in  the  same  order  as  their  atomic  weights,  while  their  melting 
points  and  boiling  points  are  in  the  inverse  order  (see  table).  Glu- 
cinum is  difficult  to  obtain  in  pure  condition,  so  that  its  constants 
have  not  been  determined  with  great  accuracy. 

At  ordinary  temperatures  oxygen  has  but  little  action  upon  the 
members  of  this  family.  At  high  temperatures,  however,  combination 
takes  place  rapidly,  with  the  formation  of  oxides,  which  have  the 
general  formula  MO,  in  which  M  represents  a  divalent  metal.  Mag- 
nesium rapidly  decomposes  boiling  water,  while  zinc  and  cadmium 
have  but  slight  action  upon  it.  They  all  dissolve  in  acids,  with  libera- 
tion of  hydrogen.  These  properties  are  in  general  accord  with  the 
position  which  the  metals  occupy  in  the  electromotive  series  (p.  158). 

Compounds.  The  members  of  the  family  are  divalent  in  then-  com- 
pounds, so  that  the  formulas  of  their  salts  resemble  those  of  the  alka- 
line earth  metals.  Like  the  latter  metals,  their  normal  carbonates, 

429 


430  GENERAL  CHEMISTEY 

phosphates,  and  silicates  are  insoluble  in  water.  Their  sulfates,  how 
ever,  are  readily  soluble.  Unlike  both  the  alkali  and  the  alkaline  earth 
metals,  the  hydroxides  of  the  metals  of  the  magnesium  family  are 
nearly  insoluble  in  water  and  are  much  more  readily  decomposed  by 
heat,  forming  water  and  the  oxide  of  the  metal.  Most  of  the  com- 
pounds ionize  in  such  a  way  as  to  give  a  simple,  colorless,  metallic  ion. 

MAGNESIUM 

Occurrence.  Magnesium  is  a  very  abundant  element  in  nature,  rank- 
ing a  little  below  calcium  in  amount.  It  never  occurs  in  an  uncom- 
bined  condition,  but  its  compounds  are  common  constituents  of  rocks 
and  are  found  in  sea  water  and  mineral  waters,  being  closely  asso- 
ciated with  the  compounds  of  calcium.  The  element  is  also  widely 
distributed  throughout  the  animal  and  vegetable  kingdoms.  In  the 
form  of  a  definite  mineral  it  occurs  as  magnesite  (MgCO3)  and 
dolomite  (CaCO3-  MgCO3)  and  is  usually  a  constituent  of  lime- 
stones. Among  its  silicates  the  following  may  be  mentioned :  asbestos 
(CaMg3(SiO3)4),  in  some  forms  of  which  the  magnesium  is  partly 
replaced  by  iron;  talc  (Mg3H2(SiO3)4) ;  serpentine  (Mg3Si2O?-  2  H2O); 
meerschaum  (Mg2H4Si3O10).  Its  presence  in  the  Stassfurt  salts  has 
already  been  noted  (p.  405). 

The  element.  Magnesium  was  first  isolated  in  1807  by  Davy,  who 
obtained  it  by  the  same  methods  which  were  successful  in  the  isolation 
of  the  alkali  and  alkaline  earth  metals.  Like  most  metals  whose  oxides 
are  difficult  to  reduce  with  carbon,  it  was  formerly  prepared  by  heating 
its  anhydrous  chloride  with  sodium : 

MgCl2  +  2  Na  =  2  NaCl  +  Mg 

At  present  it  is  prepared  by  the  electrolysis  of  anhydrous  carnallite. 
The  mineral  is  melted  in  an  iron  pot  which  serves  as  the  cathode, 
while  the  anode  is  a  carbon  rod  dipping  into  the  melted  salt. 

Magnesium  is  a  silvery  white  metal.  Its  density  is  1.74,  its  melting 
point  633°,  and  its  boiling  point  1120°.  Air  does  not  act  rapidly 
upon  it,  but  a  thin  film  of  oxide  forms  upon  its  surface,  dimming  its 
bright  luster.  It  combines  directly  with  most  of  the  nonmetals,  even 
with  nitrogen ;  hence  its  use  by  Ramsay  in  the  isolation  of  argon.  It 
is  a  strong  reducing  agent.  When  heated  in  the  air  it  is  easily  ignited 
and  burns  with  a  brilliant  white  light,  forming  the  oxide,  together 
with  a  small  percentage  of  the  nitride.  This  light  is  very  rich  in  the 


THE  MAGNESIUM  FAMILY  431 

rays  which  affect  a  photographic  plate,  so  that  the  powdered  metal, 
either  alone  or  mixed  with  potassium  chlorate,  is  used  as  a  source  of 
artificial  light  in  photography,  as  well  as  in  pyrotechnics.  It  is  also 
used  in  the  manufacture  of  magnalium,  a  light  alloy  of  magnesium 
and  aluminium  having  a  high  tensile  strength. 

Compounds  of  magnesium.  The  preparation  of  the  compounds  of 
magnesium  presents  no  new  problems.  The  carbonate  found  in  nature 
is  readily  converted  into  other  compounds  by  the  action  of  acids. 
Moreover,  magnesium  chloride  is  obtained  as  a  by-product,  especially 
in  the  preparation  of  potassium  chloride  (p.  406),  and  since  this  is  very 
soluble,  it  serves  as  a  material  for  the  preparation  of  the  insoluble  salts. 

Magnesium  oxide  (MgO);  magnesium  hydroxide  (Mg(OH)2).  Mag- 
nesium oxide  can  be  prepared  by  any  of  the  general  methods  for  pre- 
paring oxides.  Commercially  it  is  obtained  by  heating  the  carbonate, 
which  is  even  more  readily  decomposed  than  calcium  carbonate.  It  is 
a  white  powder,  very  soft  and  light,  and  is  often  known  commercially 
as  magnesia  or  magnesia  usta.  It  resembles  calcium  oxide  in  many 
respects,  but  is  even  more  infusible,  although  it  can  be  melted  in  the 
electric  furnace.  With  water  it.  forms  the  hydroxide,  but  the  combi- 
nation takes  place  much  more/  slowly,  and  with  the  production  of  less 
heat,  than  in  the  case  of  calcium  oxide.  Because  of  its  highly  infusible 
character  it  is  used  in  the  manufacture  of  fire  brick,  as  a  lining  for 
furnaces,  and  for  other  purposes  where  a  highly  refractory  substance 
is  needed. 

Magnesium  hydroxide  is  an  amorphous  substance  and  is  but  slightly 
soluble  in  water.  When  heated  it  is  easily  decomposed  into  the  oxide 
and  water.  Magnesium  salts,  as  a  rule,  have  no  injurious  effect  upon 
the  system,  and  for  this  reason  either  magnesium  oxide  or  hydroxide 
serves  as  a  very  suitable  antidote  for  poisoning  by  strong  acids,  for, 
since  they  are  basic,  they  neutralize  the  acid,  and  neither  the  excess 
taken  nor  the  salt  formed  causes  injury. 

Precipitation  of  magnesium  hydroxide.  Magnesium  hydroxide,  being  insoluble, 
is  precipitated  when  a  solution  of  sodium  or  potassium  hydroxide  is  added  to  a 
solution  of  a  magnesium  salt.  With  ammonium  hydroxide,  however,  only  a  por- 
tion of  the  magnesium  is  precipitated.  This  is  due  to  the  fact  that  ammonium 
hydroxide  is  only  slightly  ionized;  and  in  the  presence  of  the  ammonium  salts 
formed  in  the  reaction  the  extent  of  this  ionization  becomes  still  less,  because 
of  the  influence  of  the  ammonium  ions  present  in  the  solution.  As  a  result  the 
concentration  of  the  hydroxyl  ions  becomes  so  small  that  the  solubility  product 
of  magnesium  hydroxide  is  no  longer  exceeded. 


432  GENERAL  CHEMISTRY 

Magnesium  chloride  (MgCl2).  This  compound  is  found  'in  many 
natural  waters  and  salt  deposits  and  is  obtained  as  a  by-product  in 
the  manufacture  of  potassium  chloride  from  carnallite.  Under  ordi- 
nary conditions  it  crystallizes  from  solutions  as  the  hexahydrate 
MgCl2  •  6  H2O,  although  a  number  of  other  hydrates  are  known. 
When  the  hydrate  is  heated,  magnesium  oxide  is  formed: 

MgCl2  •  6  H2O  =  MgO  +  2  HC1  +  5  H2O 

Owing  to  the  abundance  of  magnesium  chloride,  attempts  have  been 
made  to  utilize  this  reaction  in  the  preparation  of  both  magesium 
oxide  and  hydrochloric  acid. 

The  anhydrous  magnesium  chloride  may  be  obtained  by  first  pre- 
paring the  double  salt  NH4C1  •  MgCl2  •  6  H2O  and  then  carefully  heat- 
ing it.  The  water  of  crystallization  is  expelled  first,  and  afterwards, 
as  the  temperature  rises,  the  ammonium  chloride  volatilizes,  leaving 
the  anhydrous  magnesium  chloride. 

Magnesium  carbonates.  The  normal  carbonate  MgCO3  occurs  in 
nature  as  magnesite  and,  combined  with  calcium  carbonate,  as  dolo- 
mite (CaCO3  •  MgCO3).  The  normal  salt  is  very  similar  to  calcium 
carbonate  in  its  properties,  but  is  slightly  more  soluble  in  water  and 
much  more  readily  decomposed  by  heat.  It  is  quite  soluble  in  water 
containing  carbon  dioxide,  owing  to  the  formation  of  the  soluble  acid 
carbonate  Mg(HCO3)2.  When  a  solution  of  magnesium  salt  is  precipi- 
tated with  sodium  or  potassium  carbonate,  a  white  solid  is  obtained 
which  is  not  the  normal  carbonate  that  one  would  naturally  expect, 
but  a  basic  salt,  the  exact  composition  of  which  varies  with  the  con- 
ditions of  the  experiment.  As  ordinarily  prepared  its  composition  is 
expressed  by  the  formula  3  MgCO8  •  Mg(OH)2  •  3  H2O.  This  basic  salt 
is  known  as  magnesia  alba  and  is  used  in  medicine  and  as  a  cosmetic. 

Magnesium  sulfate  (MgSOJ.  Like  the  chloride,  this  salt  is  found 
in  many  salt  beds,  and  deposits  of  the  nearly  pure  compound  occur  in 
Wyoming  and  Washington.  A  number  of  hydrates  are  known.  The 
monohydrate  occurs  in  the  Stassfurt  deposits  as  kieserite,  but  the 
most  common  form  is  the  heptahydrate  MgSO4  •  7  H2O,  known  as 
Epsom  salts.  This  form  is  obtained  by  crystallizing  magnesium  sul- 
fate from  solution  at  ordinary  temperature,  and,  unlike  kieserite,  it 
is  very  soluble  in  water.  All  the  hydrates  yield  the  anhydrous  salt 
when  heated  to  200°,  and  this  in  turn  decomposes  at  a  white  heat, 
leaving  a  residue  of  magnesium  oxide. 


THE  MAGNESIUM  FAMILY  433 

Magnesium  sulfate  has  many  uses  in  the  industries.  It  is  used  as 
a  coating  for  cotton  cloth,  in  the  dyeing  industry,  in  tanning,  in  the 
manufacture  of  paints  and  laundry  soaps,  and  to  a  limited  extent  in 
the  preparation  of  sodium  and  potassium  sulfates  and  in  medicine. 

Magnesium  sulfate  was  the  first  magnesium  compound  to  be  described.  In 
1695  Grew,  a  London  physician,  called  attention  to  its  occurrence  in  the  waters 
of  the  famous  spring  located  at  Epsom,  England ;  hence  the  name  Epsom  salts. 
It  soon  came  into  general  use  as  a  medicinal  agent. 

Boiler  scale.  When  water  which  contains  certain  salts  in  solution  is  evaporated 
in  steam  boilers,  a  hard,  insoluble  material,  called  scale,  deposits  in  the  boiler. 
The  formation  of  this  scale  may  be  due  to  several  distinct  causes. 

1.  To  the  deposit  of  calcium  sulfate.  This  salt,  while  sparingly  soluble  in  cold 
water,  is  almost  completely  insoluble  in  superheated  water.    Consequently,  it  is 
precipitated  when  water  containing  it  is  heated  in  a  boiler. 

2.  To  decomposition  of  acid  carbonates.  As  we  have  seen,  calcium  and  magnesium 
acid  carbonates  are  decomposed  on  heating,  forming  insoluble  normal  carbonates  : 

Ca(HCO3)2  =  CaCO3  +  H2O  +  CO2 

3.  To  hydrolysis  of  magnesium  salts.  Magnesium  chloride  and  to  some  extent 
magnesium  sulfate  undergo  hydrolysis  when  superheated  in  solution,  and  the 
magnesium  hydroxide,  being  sparingly  soluble,  precipitates : 

MgCl2  +  2  H20^=>:Mg(OH)2  +  2  HC1 

This  scale  adheres  tightly  to  the  boiler  in  compact  layers  and,  being  a  noncon- 
ductor of  heat,  causes  much  waste  of  fuel.  It  is  very  difficult  to  remove,  owing 
to  its  hardness  and  resistance  to  reagents.  Thick  scale  sometimes  cracks,  and 
the  water,  coming  in  contact  with  the  overheated  iron,  may  cause  an  explosion. 
Moreover,  the  acids  set  free  in  the  hydrolysis  of  the  magnesium  salts  attack  the 
iron  tubes  and  rapidly  corrode  them.  These  causes  combine  to  make  the  forma- 
tion of  scale  a  matter  which  causes  much  trouble  in  cases  where  hard  water  is 
used  in  steam  boilers.  Water  containing  such  salts  should  be  softened,  therefore, 
before  being  used  in  boilers. 

Other  magnesium  compounds.  Magnesium  sulfide  (MgS)  is  prepared 
by  heating  a  mixture  of  magnesium  and  sulfur.  It  is  a  yellowish- 
gray  solid  and  resembles  calcium  sulfide  in  its  action  toward  water. 
Magnesium  nitride  (MggN2)  is  formed  by  the  direct  union  of  magnesium 
and  nitrogen  at  high  temperature.  It  reacts  with  water,  forming  mag- 
nesium hydroxide  and  ammonia.  The  phosphates  of  magnesium  resemble 
those  of  calcium  in  their  composition  and  general  properties.  When 
a  solution  of  disodium  phosphate  is  added  to  a  'solution  of  any 
magnesium  compound  containing  ammonium  hydroxide,  magnesium 
ammonium  phosphate  is  precipitated: 

MgS04  +  Na2HP04  +  NH4OH  =  MgNH4PO4  +  Na2SO4  +  H2O 


434  GENERAL  CHEMISTRY 

This  crystallizes  in  the  form  of  the  hexahydrate,  which  is  somewhat 
unstable.  If  filtered  off  and  heated  to  redness,  however,  it  is  converted 
into  magnesium  pyrophosphate  (Mg2P2O7),  which  is  perfectly  stable : 

2  MgNH4PO4  =  Mg2P2O7  +  2  NH3  +  H2O 

From  the  weight  of  the  pyrophosphate  it  is  possible  to  calculate  the 
weight  of  magnesium  present  in  the  original  solution.  These  reactions 
serve  in  the  laboratory  for  the  quantitative  determination  either  of 
magnesium  or  of  phosphoric  acid. 

ZINC 

Occurrence  and  metallurgy.  Zinc  does  not  occur  free  in  nature.  It 
is  not  a  constituent  of  common  rocks,  and  its  occurrence  is  rather 
local  and  confined  to  deposits  or  pockets.  Its  chief  ores  are  the  fol- 
lowing :  sphalerite  (ZnS) ;  zincite  (ZnO) ;  smithsonite  (ZnCO3) ; 
willemite  (Zn2SiO4) ;  calamine  (Zn2SiO4  •  H2O) ;  franklinite  (ZnFe2O4). 
In  the  United  States  it  occurs  most  abundantly  in  Kansas,  Missouri, 
and  New  Jersey. 

In  the  metallurgy  of  zinc  the  ores  are  first  converted  into  the 
oxide  by  roasting,  and  the  oxide  is  then  reduced  with  carbon. 

A  mixture  of  zinc  oxide  and  coal  is  heated  in  earthenware  retorts.  The  zinc 
oxide  is  thereby  reduced,  and  the  resulting  zinc,  being  volatile  at  the  temperature 
of  the  retort,  distills  and  is  collected  in  suitable  receivers.  At  first  the  zinc  col- 
lects in  the  form  of  a  powder  known  as  zinc  dust,  which,  however,  contains  some 
zinc  oxide.  Later,  when  the  receiver  has  become  hot,  the  zinc  condenses  to  a 
liquid  and  is  drawn  off  into  molds.  In  this  form  it  is  impure  and  is  known 
as  spelter. 

Commercial  zinc  generally  contains  a  number  of  impurities,  espe- 
cially carbon,  arsenic,  cadmium,  and  iron.  These  can  be  largely  re- 
moved by  distillation.  Zinc  containing  less  than  0.001  per  cent  of 
impurities  has  been  obtained  by  electrolysis  of  a  pure  salt  and  sub- 
sequent distillation  under  diminished  pressure. 

Propertied.  Pure  zinc  is  a  bluish-white  metal.  Its  density  is  7.10, 
its  melting  point  419.4°,  and  its  boiling  point  918°.  Some  of  its  physi- 
cal properties  are  greatly  modified  by  the  temperature  and  the  previ- 
ous treatment  of  the  metal.  When  allowed  to  solidify  from  the  liquid 
state,  it  is  highly  crystalline  and  is  quite  hard  and  brittle.  At  tem- 
peratures between  100°  and  150°  it  is  malleable  and  can  be  rolled 
into  thin  sheets,  which  retain  their  softness  and  malleability  at  ordinary 


THE  MAGNESIUM  FAMILY  435 

temperatures.  Above  150°  it  again  becomes  very  brittle.  When  melted 
and  poured  slowly  into  water,  it  forms  thin,  brittle  flakes,  and  in  this 
condition  is  called  granulated  zinc  or  mossy  zinc. 

Zinc  tarnishes  superficially  in  moist  air,  owing  to  the  formation  of  a 
basic  carbonate.  It  does  not  decompose  even  boiling  water,  but  at  a 
high  temperature  it  acts  upon  steam,  forming  the  oxide  and  hydrogen. 
When  heated  sufficiently  in  oxygen  or  air  it  burns  with  a  bluish  flame. 
Dilute  acids  have  but  little  action  upon  the  pure  metal,  since  the  hydro- 
gen at  first  liberated  collects  on  the  surface  of  the  metal  in  the  form 
of  a  thin  film,  and  the  action  soon  ceases.  If  another  metal  below  zinc 
in  the  electromotive  series,  such  as  iron,  copper,  or  platinum,  is  present, 
either  as  an  impurity  or  simply  placed  in  contact  with  the  zinc,  the 
hydrogen  escapes  from  the  surface  of  the  metal  having  the  lower 
electrode  potential,  and  the  zinc  rapidly  dissolves.  Zinc  also  dissolves 
in  sodium  and  potassium  hydroxides,  with  liberation  of  hydrogen: 

Zn  +  2  KOH  =  Zn(OK)2  •+  H2 

Uses  of  zinc.  The  chief  use  of  zinc  is  in  the  manufacture  of  gal- 
vanized iron.  This  is  sheet  iron  covered  with  a  thin  layer  of  zinc, 
which  protects  the  iron  from  the  action  of  air  and  water.  About,  two 
thirds  of  all  the  zinc  produced  is  used  in  this  way.  Large  quantities 
are  also  used  in  the  manufacture  of  alloys  (table,  p.  480).  The  metal 
is  also  used  in  the  construction  of  batteries  and  as  a  roofing  material. 
In  the  laboratory  it  is  used  in  the  preparation  of  hydrogen  and,  in  the 
form  of  zinc  dust,  as  a  reducing  agent. 

Compounds  of  zinc.  In  general  the  compounds  of  zinc  are  similar  in 
formula  and  appearance  to  those  of  magnesium.  They  often  differ  from 
them  quite  markedly,  however,  in  chemical  conduct.  Either  the  metal, 
its  oxide,  or  the  natural  carbonate  or  sulfide  serves  as  a  convenient 
material  for  the  preparation  of  these  compounds. 

Zinc  oxide  (zinc  white)  (ZnO).  This  is  a  white  powder  obtained  by 
roasting  the  ores  in  a  current  of  air  or,  in  the  pure  state,  by  oxidizing 
the  metal  itself.  It  turns  yellow  when  heated,  but  regains  its  white  color 
on  cooling.  It  is  used  very  largely  as  a  white  pigment  in  paints,  under 
the  name  of  zinc  white,  and  has  an  advantage  over  white  lead  in  that 
it  is  not  darkened  by  the  sulfur  compounds  which  are  present  in  the 
air,  especially  in  manufacturing  districts.  It  is  said  that  at  present 
40,000,000  Ib.  of  zinc  oxide  are  used  annually  in  the  manufacture 
of  rubber\gopds,  especially  as  a  filler  for  automobile  tires. 


•^-rn  ^>  o^  3  (-?/  -       w-  i 

436  GENERAL  CHEMISTRY 

Zinc  hydroxide  (Zn(OH)2).  This  compound  is  precipitated  in  the 
form  of  a  white  solid  when  an  alkali  is  added  to  a  solution  of  a  zinc 
salt.  It  is  an  amphoteric  hydroxide,  dissolving  in  both  acids  and 
alkalies.  In  solution  it  ionizes  as  follows: 

H+,  H+,  ZnO2--^±Zn(OH)2^=>:Zn++,  OH~,  OH~ 

Soluble  salts.  The  soluble  salts  of  zinc  can  be  prepared  by  dissolv- 
ing the  metal  or  its  oxide  in  the  appropriate  acid.  The  chloride  and 
sulfate  are  the  most  familiar. 

Zinc  chloride  (ZnCl2).  This  salt  is  readily  soluble  in  water  and  in 
alcohol.  It  is  strongly  hydrolyzed  in  water,  and  upon  evaporation 
the  solution  yields  a  basic  chloride  ZnOHCl  in  addition  to  zinc  chloride. 
When  fused  and  allowed  to  cool,  it  forms  a  hard  mass  which  has  a 
marked  affinity  for  water  and  is  strongly  caustic.  The  largest  use  of 
zinc  chloride  is  as  a  wood  preservative. 

Zinc  sulfate  (ZnSOJ.  Under  ordinary  conditions  this  salt  crystal- 
lizes from  water  in  the  form  of  a  heptahydrate  ZnSO4  •  7  H2O,  which 
has  long  been  known  under  the  name  of  white  vitriol.  Commercially 
the  salt  is  prepared  by  roasting  sphalerite.  It  is  used  in  medicine,  and 
to  a  limited  extent  in  the  dyeing  and  printing  of  cloth. 

Insoluble  salts.  These  are  prepared  by  the  ordinary  methods  of 
precipitation.  The  most  important  ones  are  the  sulfide  and  carbonate. 
The  normal  carbonate  ZnCOg  is  precipitated  upon  the  addition  of 
sodium  bicarbonate  to  a  solution  of  a  zinc  salt.  Normal  sodium  car- 
bonate precipitates  basic  salts,  as  in  the  case  of  magnesium. 

Zinc  sulfide  (ZnS).  The  natural  zinc  sulfide  (sphalerite,  blende,  or 
blackjack)  varies,  largely  according  to  the  impurities  present,  from  a 
light-yellow  transparent,  resinoufs  solid  to  a  black  mass.  As  prepared 
in  the  laboratory  by  precipitation,  it  is  a  white  solid.  It  is  important 
to  note  that  zinc  is  the  only  one  of  the  common  metals  that  forms  a 
white  sulfide  nearly  insoluble  in  water. 

Formation  of  zinc  sulfide  by  precipitation.  Zinc  sulfide  is  nearly 
insoluble  in  water  and  is  formed  as  a  precipitate  when  ammonium 
sulfide  is  added  to  a  solution  of  a  salt  of  zinc: 

ZnCl2  +  (NH4)2S  =  ZnS  +  2  NH4C1 

On  the  other  hand,  when  hydrogen  sulfide  is  passed  into  such  a  solu- 
tion, the  precipitation  of  the  zinc  sulfide  soon  ceases,  an  equilibrium 
resulting,  as  expressed  in  the  following  equation : 
ZnCl2  +  H2S  +=±  ZnS  +  2  HC1 


THE  MAGNESIUM  FAMILY  437 

In  accordance  with  the  theory  of  precipitation  (p.  388)  this  equilibrium  results 
in  the  following  way  :  While  zinc  sulfide  is  commonly  said  to  be  insoluble,  because 
it  is  nearly  so,  it  really  has  a  very  perceptible  solubility.  Before  it  can  begin  to 
precipitate,  the  concentration  of  the  zinc  ions  and  sulfur  ions  must  be  large 
enough  to  exceed  the  solubility  product  of  the  zinc  sulfide,  as  expressed  in  the 
following  equation  : 


There  is  an  abundant  concentration  of  the  zinc  ions  supplied  by  the  zinc  salt; 
moreover,  at  the  beginning  of  the  reaction  the  concentration  of  the  sulfur  ions 
derived  from  the  hydrogen  sulfide  is  fairly  large,  so  that  the  solubility  product  of 
the  zinc  sulfide  is  exceeded  and  precipitation  takes  place.  As  the  reaction  pro- 
ceeds, however,  increasing  quantities  of  hydrogen  chloride  are  formed,  and  the 
hydrogen  ions  derived  from  this  compound  diminish  the  extent  of  the  ionization 
of  the  hydrogen  sulfide.  The  concentration  of  the  sulfur  ions,  therefore,  gradu- 
ally decreases  until  the  .solubility  product  of  the  zinc  sulfide  is  no  longer  exceeded 
and  precipitation  ceases.  It  is  interesting  to  note,  however,  that  if  a  normal  salt 
of  some  weak  acid,  such  as  sodium.  acetate  (NaC2H3O2)>  is  added  to  the  solution 
of  the  -zinc  salt,  and  the  hydrogen  sulfide  then  passed  in,  the  precipitation  of  the 
zinc  sulfide  continues  to  completion,  for  as  fast  as  the  hydrogen  ions  are  formed 
in  the  reaction,  they  enter  into  an  equilibrium  with  the  anion  of  the  weak  acid, 
as  shown  in  the  following  equation  : 

H+  +  C2H302-^=±HC2H302 

This  keeps  the  concentration  of  the  hydrogen  ions  too  low  to  have  any  marked 
effect  upon  the  concentration  of  the  sulfur  ions,  so  that  precipitation  continues 
to  completion.  It  may  be  added  that  the  addition  of  sodium  acetate  to  an  acid 
solution  is  a  device  often  employed  when  it  is  desired  to  reduce  the  concentration 
of  the  hydrogen  ions  to  a  minimum  value  and  yet  have  the  solution  distinctly 
acid  in  reaction. 

CADMIUM 

Preparation  and  properties.  This  metal  is  associated  with  zinc  in 
nature,  small  quantities  occurring  in  many  zinc  ores.  In  the  course 
of  the  metallurgy  of  zinc  the  cadmium  compounds  are  changed  to 
cadmium  oxide,  which  is  then  reduced  by  carbon,  the  cadmium  dis- 
tilling over  with  the  zinc.  Being  more  volatile  than  the  zinc,  it  is 
largely  concentrated  in  the  first  portions  of  the  distillate.  From  these 
portions  the  pure  metal  can  be  obtained  by  fractional  distillation.  The 
metal  resembles  zinc  in  appearance.  Its  density  is  8.64,  its  melting 
point  321°,  and  its  boiling  point  778°. 

It  is  used  to  a  limited  extent  in  making  colored  glass  and  alloys. 
It  is  also  used  to  some  extent  in  the  construction  of  standard  cells 
(p.  490).  Its  alloys  are  in  general  characterized  by  a  low  melting 
point  (p.  375). 


438  GENERAL  CHEMISTRY 

Only  small  quantities  of  cadmium  are  produced,  since  there  is  little  demand  for 
it  at  present  prices.  The  zinc  spelter  obtained  from  the  ore  of  the  Western  states 
contains  about  0.4  per  cent  of  cadmium.  Although  its  presence  in  the  zinc  is 
undesirable,  yet  under  present  conditions  it  is  not  economical  to  separate  it. 

Compounds  of  cadmium.  Some  of  the  most  important  of  the  com- 
pounds of  cadmium  are  the  following:  Cadmium  oxide  (CdO)  is 
obtained  by  heating  cadmium  in  air  or  oxygen.  It  is  a  brown  powder 
and  combines  with  water  to  form  cadmium  hydroxide  (Cd(OH)2),  a 
white,  insoluble  solid.  Cadmium  chloride  (CdCl2)  crystallizes  from 
water  as  the  dihydrate.  It  does  not  hydrolyze  as  does  zinc  chloride. 
Cadmium  bromide  (CdBr2)  and  cadmium  iodide  (CdI2)  resemble  the 
chloride  and  are  used  to  some  extent  in  photography.  Cadmium  sulfate 
crystallizes  from  water  as  a  hydrate  of  the  formula  3  CdSO4  •  8  H2O, 
while  under  similar  conditions  cadmium  nitrate  forms  the  hydrate 
Cd(NO3)2  •  4  H2O.  Cadmium  sulfide  (CdS)  is  commercially  the  most 
important  of  the  cadmium  compounds.  It  occurs  in  nature  in  the 
form  of  the  mineral  known  as  greenockite.  In  the  laboratory  it  is 
prepared  by  passing  hydrogen  sulfide  into  a  solution  of  a  cadmium 
salt.  It  is  bright  yellow  in  color  and  is  used  as  a  pigment. 

GLUCINUM 

Preparation  and  properties.  Glucinum,  known  also  as  beryllium, 
since  it  was  first  found  in  the  mineral  beryl,  is  a  rather  rare  metal. 
In  1828  Wohler  first  isolated  it  by  heating  the  chloride  with  potas- 
sium. It  is  a  hard  metal  and  has  a  bright,  metallic  luster.  When  pre- 
pared by  electrolysis,  crystals  are  obtained  which  have  a  density  of 
1.93.  It  is  similar  to  the  other  metals  of  the  family  in  its  behavior 
toward  air  and  acids.  It  resembles  zinc  in  that  it  readily  dissolves 
in  alkalies. 

Compounds  of  glucinum.  This  element  forms  a  series  of  compounds 
resembling  in  formulas  the  corresponding  compounds  of  the  other  mem- 
bers of  the  magnesium  family.  The  hydroxide  G1(OH)2,  like  zinc 
hydroxide,  is  insoluble  in  water  but  dissolves  in  both  acids  and  alkalies. 
The  salts  of  glucinum  readily  hydrolyze  and  form  basic  compounds. 
The  soluble  compounds  of  the  element  have  a  sweetish  taste,  which  fact 
suggested  the  name  glucinum,  from  the  Greek  word  meaning  "  sweet." 


CHAPTER  XXXII 

THE  ALUMINIUM  GROUP 

The  group.  With  the  exception  of  aluminium,  none  of  the  ele- 
ments of  Group  III  of  the  periodic  table  are  well  known  or  abundant. 
Boron  has  already  been  considered,  and  the  others  fall  naturally  into 
two  families.  The  one  includes  aluminium,  together  with  gallium, 
indium,  and  thallium;  the  other,  scandium  and  yttrium,  together 
with  a  large  group  of  elements  whose  oxides  are  collectively  called 
the  rare  earths. 

All  of  the'  elements  of  this  group  are  trivalent  in  their  compounds, 
though  some  of  the  rarer  elements,  particularly  thallium,  have  lower 
valences  as  well.  With  few  exceptions  their  salts  are  colorless,  save 
when  they  are  derived  from  a  colored  acid.  The  bases  which  these 
elements  form  are  nearly  all  quite  weak,  and  many  of  their,  salts  are 
hydrolyzed  in  solution.  A  brief  mention  of  the  rarer  elements  will 
be  made  after  aluminium  has  been  considered. 

ALUMINIUM 

Occurrence.  Next  to  oxygen  and  silicon,  aluminium  is  the  most 
abundant  of  all  the  elements.  The  free  element  is  not  found  in  nature, 
but  its  compounds,  especially  the  silicates,  are  abundant  and  widely 
distributed,  being  essential  constituents  of  all  important  soils  and 
rocks  excepting  limestone  and  sandstone.  The  feldspars,  which  are 
the  most  abundant  of  all  the  minerals  in  the  earth's  crust,  are  all  sili- 
cates of  aluminium  and  either  sodium,  potassium,  or  calcium.  Since 
the  soil  has  been  formed  largely  by  the  disintegration  of  these  rocks, 
it  is  rich  in  the  silicates  of  aluminium,  chiefly  in  the  form  of  clay. 
Some  of  the  other  forms  in  which  aluminium  occurs  in  nature  are  the 
following :  corundum  (A12O3') ;  emery  (A12O3  colored  black  with  oxide 
of  iron);  cryolite  (Na3AlFg);  bauxite,  a  mixture  of  iron  oxicje  and 
hydrated  aluminium  oxides  (A12O8  •  H2O  and  A12O3  •  3  H2O).  Bauxite 
is  the  ore  from  which  aluminium  is  prepared.  In  the  United  States  it 
is  found  chiefly  in  Georgia,  Alabama,  and  Arkansas. 

439 


440 


GENERAL   CHEMISTRY 


Preparation.  Aluminium  was  first  prepared  by  Wohler  in  1827 
by  heating  anhydrous  aluminium  chloride  with  potassium: 

A1C13  +  3  K  =  3  KC1  +  Al 

Although  the  metal  is  very  abundant  in  nature,  and  possesses  many 
desirable  properties,  the  cost  of  separating  it  from  its  ores  by  the 
earlier  methods  was  so  great  that  it  remained  almost  a  curiosity  until 
comparatively  recent  years,  when  greatly  improved  methods  of  prepa- 
ration came  into  use.  It  is  now  prepared  by  the  electrolysis  of  alu. 
minium  oxide  (A12O3)  dissolved  in  melted  cryolite  —  a  method  first 
patented  by  the  American  chemist  Hall  in  1886.  As  a  result  of  this 
improvement  the  production  of  aluminium  increased  from  83  Ib.  in 
1883  to  nearly  50,000,000  Ib.  in  1911,  while  the  price  of  the  metal 
has  decreased  per  pound  during  the  last  twenty  years  from  $5.00 
to  $0.20. 

The  commercial  preparation  of  aluminium.  An  iron  box  A  (Fig.  139)  about 
eight  feet  long  and  six  feet  wide  is  connected  with  a  powerful  electrical  generator 
in  such  a  way  as  to  constitute  the  cathode  upon  which  the  aluminium  is  deposited. 
Three  or  four  rows  of  carbon  rods  B  dip  into  the  box  and  serve  as  the  anodes. 
The  box  is  partially  filled  with  cryolite  and  the  current  is  turned  on,  generating 


FIG.  139 


sufficient  heat  to  melt  the  cryolite.  Aluminium  oxide  is  then  added,  and  under 
the  influence  of  the  electric  current  it  decomposes  into  aluminium  and  oxygen.  - 
The  temperature  is  maintained  above  the  melting  point  of  aluminium,  and  the 
liquid  metal,  being  heavier  than  cryolite,  sinks  to  the  bottom  of  the  vessel,  from 
which  it  is  tapped  off  from  time  to  time  through  the  tap  hole  C.  The  oxygen  in 
part  escapes  as  gas,  and  in  part  combines  with  the  carbon  of  the  anode. 

The  largest  expense  in  the  process,  apart  from  the  cost  of  electrical  energy, 
is  the  preparation  of  aluminium  oxide  free  from  other  oxides,  for  most  of  the 


THE  ALUMINIUM  GROUP  441 

oxide  found  in  nature  'is  too  impure  to  serve  without  refining.  Bauxite  is  used 
as  the  ore,  because  it  is  converted  into  pure  oxide  without  great  difficulty. 
Since  common  clay  is  a  silicate  of  aluminium  and  is  everywhere  abundant,  it 
might  be  expected  that  this  would  be  utilized  in  the  preparation  of  aluminium. 
It  is,  however,  very  difficult  to  extract  the  aluminium  from  a  silicate,  and  no 
practical  method  has  been  found  which  will  accomplish  this. 

Properties.  Aluminium  resembles  tin  in  appearance.  Its  density  is 
2.65,  being  only  about  one  third  that  of  iron.  It  melts  at  658.5°.  It 
is  ductile  and  malleable,  especially  at  temperatures  between  100°  and 
150°,  when  it  can  be  hammered  into  very  thin  sheets.  At  higher  tem- 
peratures, near  its  melting  point,  it  is  very  brittle.  It  is  fairly  hard 
and  strong,  being  superior  to  most  metals  in  these  respects,  although 
not  equal  to  steel.  It  is  an  excellent  conductor  of  heat  and  electricity. 

Aluminium  is  but  slightly  acted  upon  by  water,  while  moist  air 
merely  dims  its  luster.  Further  action  is  prevented  in  each  case  by 
the  formation  of  a  very  thin  film  of  oxide  upon  the  surface  of  the 
metal.  It  combines  with  many  of  the  nonmetals,  especially  with  the 
halogens  and  the  members  of  the  sulfur  family.  It  is  an  excellent 
reducing  agent,  combining  with  oxygen  at  high  temperatures,  with 
liberation  of  much  heat: 

4  Al  +  3  02  =  2  A1203  +  760,400  cal. 

Nitric  acid  and  dilute  sulfuric  acid  have  but  little  action  upon  it; 
concentrated  sulfuric  acid  dissolves  it,  forming  the  sulfate  and  liber- 
ating sulfur  dioxide.  Hydrochloric  acid  is  its  best  solvent: 

2  A1  +  6  HC1  =  2  A1C1,  -h  3  H2 

Aluminium  resembles  zinc  in  that  it  readily  dissolves  in  strong  alkalies, 
forming  aluminates  and  liberating  hydrogen.  It  is  also  acted  upon  by 
sodium  chloride,  especially  in  the  presence  of  oxygen  and  dilute  acids 
such  as  acetic. 

Uses.  The  lightness  and  strength  of  aluminium,  together  with  its 
inertness  toward  air  and  water,  suggest  a  variety  of  applications  for  the 
metal.  It  is  used  for  many  construction  purposes  and  for  the  manu- 
facture of  cooking  utensils.  In  the  form  of  a  powder  suspended  in  a 
suitable  liquid,  it  makes  an  efficient  silver-like  paint.  Although  not 
so  good  a  conductor  of  electricity  as  copper  for  a  given  cross  section 
of  wire,  nevertheless,  weight  for  weight,  it  is  an  even  better  conductor 
and  is  coming  into  use  in  electrical  construction,  especially  for  long- 
distance power  wires.  It  is  also  used  to  quite  an  extent  as  a  reducing 


442  GENERAL  CHEMISTRY 

agent  in  the  Goldschmidt  process  (p.  385).  The  greatest  use  of 
aluminium,  however,  is  in  the  iron  and  steel  industries  and  in  the 
manufacture  of  alloys.  A  small  quantity  is  often  added  to  molten 
steel  in  order  to  combine  with  any  oxygen  present  and  thus  prevent 
the  formation  of  bubbles  and  cavities  in  the  metal.  Aluminium  bronze, 
consisting  of  about  90  per  cent  copper  and  10  per  cent  aluminium,  has 
a  pure  golden  color,  is  strong  and  malleable,  is  easily  cast,  and  is  per- 
manent in  the  air.  Magnalium  (p.  431)  is  silver-white  and  very  light. 

Goldschmidt  welding  process.  The  property  possessed  by  aluminium  of  reducing 
oxides  with  the  liberation  of  a  large  amount  of  heat  is  turned  into  practical 
account  in  the  welding  of  metals.  The  German  chemist  Goldschmidt  was  the 
first  to  use  aluminium  for  this  purpose.  The  welding 
of  metals  by  this  method  may  be  illustrated  by  a  single 
example,  namely,  the  welding  of  car  rails  —  a  process 
often  carried  out  in  connection  with  electric  railways 
to  secure  good  electrical  connection.  The  ends  of  the 
rails  are  accurately  aligned  and  thoroughly  cleaned. 
A  sand  mold  A  (Fig.  140)  is  then  clamped  about  the 
ends  of  the  rails,  leaving  sufficient  space  so  that  the 
metal  can  flow  in.  The  ends-  of  the  rails  are  heated  to 
redness  by  the  flame  from  a  gasoline  compressed-air 
torch  directed  into  the  opening  in  the  mold.  Just  over 
the  opening  is  placed  the  conical-shaped  crucible  B, 
which  contains  a  mixture  of  iron,  metallic  oxides,  and 
aluminium.  When  the  ends  of  the  rails  have  been 
FIG  140  heated  to  redness  by  the  torch,  the  mixture  in  the 

crucible  is  ignited,  and  after  a  few  seconds  the  crucible 

is  opened  at  the  bottom,  and  the  molten  metal  resulting  from  the  reaction  in  the 
crucible  is  allowed  to  flow  into  the  mold.  In  this  way  the  molten  metal  sur- 
rounds the  ends  of  the  rails  and,  as  it  cools,  welds  them  firmly  together.  A  mix- 
ture of  the  metallic  oxides  and  aluminium  ready  for  use  in  welding  is  sold  under 
the  name  of  thermite. 

Compounds  of  aluminium.  Aluminium  is  a  trivalent  metal,  and  the 
formulas  of  its  compounds  therefore  resemble  those  of  bismuth  and 
antimony.  Aluminium  hydroxide,  like  antimony  hydroxide,  is  am- 
photeric.  With  strong  bases  it  forms  aluminates  such  as  sodium 
aluminate  Al(ONa)3  or  NagAlO3,  while  with  acids  it  forms  salts  such 
as  the  chloride  A1C18  and  the  sulfate  A12(SO4)3.  These  salts  are  char- 
acterized by  their  great  tendency  to  undergo  hydrolysis.  Aqueous 
solutions  of  the  chloride  and  sulfate  are  strongly  acid  in  reaction, 
while  the  carbonate  and  sulfide  are  completely  decomposed  by 
water  (p.  226). 


THE  ALUMINIUM  GROUP  443 

Aluminium  oxide  (A1203).  The  occurrence  of  aluminium  in  nature 
in  the  form  of  corundum  and  emery  has  already  been  mentioned.  In 
transparent  crystals  tinted  different  colors  by  traces  of  other  substances, 
such  as  manganese  and  chromium,  it  forms  such  precious  stones  as  the 
sapphire,  ruby,  oriental  amethyst,  and  oriental  topaz.  All  of  these 
are  very  hard,  falling  but  little  short  of  the  diamond  in  this  respect. 
The  cheaper  forms  (corundum  and  emery)  are  therefore  used  as  abra- 
sives. By  igniting  the  hydroxide  the  pure  aluminium  oxide  may  be 
obtained  in  the  form  of  a  white,  amorphous  powder: 

2A1(OH)3  =  A1203  +  3H20 

When  heated  to  about  1900°  it  melts,  and  on  cooling  forms  a  crys- 
talline mass  resembling  natural  corundum.  Some  forms  of  laboratory 
apparatus,  such  as  crucibles  and  tubes,  are  being  made  of  aluminium 
oxide.  When  used  for  this  purpose  the  oxide  is  known  as  alundum. 

Laboratory  preparation  of  gems.  A  number  of  gems  are  now  prepared  in  the 
laboratory  from  molten  aluminium  oxide.  The  white  sapphires  so  extensively 
advertised  are  simply  the  pure  oxide.  By  incorporating  with  the  melted  oxide 
small  percentages  of  certain  metallic  oxides,  different  tints  or  colors  are  obtained, 
and  in  this  way  are  prepared  such  gems  as  the  ruby,  the  oriental  amethyst,  and 
the  yellow  and  blue  sapphires,  which  are  practically  identical  in  composition 
and  properties  with  the  natural  stones. 

Aluminium  hydroxide  (A1(OH)3)  or  aluminic  acid  (H3A103).  This 
compound  may  be  prepared  by  adding  ammonium  hydroxide  to  a 
solution  of  an  aluminium  salt  : 


A1C13  +  3  NH4OH  =  3  NH4C1  +  A1(OH)8 

It  forms  a  colloidal  solution  from  which  the  insoluble  hydrogel  separates 
and  slowly  settles  in  the  liquid  in  the  form  of  a  white,  gelatinous  solid. 
It  is  amphoteric  in  character,  ionizing  as  represented  in  the  following 
equation  :  AJ+  +  +?  3  QH_  _+.  A1(OH)3  •<->•  3  H+,  A1O$-  - 

When  treated  with  a  concentrated  solution  of  sodium  hydroxide,  it 
dissolves,  owing  to  the  formation  of  a  soluble  aluminate.  This  is  diffi- 
cult to  obtain  in  pure  form,  and  its  composition  has  not  been  definitely 
settled.  One  would  naturally  expect  the  reaction  to  take  place  as 
indicated  in  the  following  equation  : 

H3A108  +  3  NaOH  =  Na8AlO3  +  3  H2O 

A  number  of  minerals,  such  as  spinel  (Mg(AlO2)2),  are  found  in  nature 
which  are  apparently  salts  of  an  acid  having  the  formula  HA1O2  and 


444  GEKEKAL   CHEMISTRY 

known  as  metaluminic  acid.  This  may  be  regarded  as  derived  from 
aluminium  hydroxide  by  the  loss  of  water,  just  as  boric  acid  yields 
metaboric  acid :  Al(OH)g  =  HA1O2  +  H2O 

Spinel  minerals.  A  number  of  other  trivalent  hydroxides  act  in  this  same  way, 
yielding  salts  analogous  to  spinel.  Many  of  these  are  important  minerals  and 
are  sometimes  called  collectively  the  spinels.  A  few  of  them  are  as  follows : 

Spinel  (Mg(AlO2)2  or  MgO  •  A12O3)  Magnetite  (Fe(FeO2)2  or  FeO  •  Fe2O3) 

Franklinite  (Zn(FeO2)2  or  ZnO  •  Fe2O3)     Chromite  (Fe(CrO2)2   or  FeO  •  Cr2O3) 

Use  of  aluminium  hydroxide.  Aluminium  hydroxide  either  combines  with 
or  absorbs  many  soluble  coloring  substances,  forming  insoluble  products.  This 
property  leads  to  its  wide  use  in  the  dyeing  industry.  Many  dyes  will  not  adhere  to 
natural  fibers,  such  as  cotton ;  that  is,  they  will  not  dye  fast.  It  is  often  possible 
to  dye  such  cloth  in  the  following  way :  The  cloth  is  first  soaked  in  a  solution 
of  an  aluminium  salt,  such  as  the  acetate,  which  readily  undergoes  hydrolysis. 
The  cloth  is  then  exposed  to  the  action  of  steam,  whereby  the  aluminium  salt  is 
completely  hydrolyzed,  the  resulting  aluminium  hydroxide  being  thus  thoroughly 
incorporated  in  the  fiber.  If  the  cloth  is  now  dipped  into  a  solution  of  the  dye, 
the  aluminium  hydroxide  combines  with  or  absorbs  the  color  substance  and 
fastens,  or  "  fixes,"  it  upon  the  fiber.  A  substance  such  as  aluminium  hydroxide 
which  serves  this  purpose  is  known  as  a  mordant,  which  means  "  biting,"  since  it 
bites,  or  holds  fast,  the  dye. 

The  value  of  aluminium  hydroxide  in  the  purification  of  water  is  due  largely 
to  its  gelatinous  character  when  freshly  formed  by  precipitation.  When  stirred 
through  the  water  it  slowly  settles,  and  in  doing  so  carries  with  it  any  sus- 
pended matter  present,  including  microorganisms.  Any  coloring  matter  present 
in  the  water  is  likewise  removed.  Instead  of  adding  the  aluminium  hydroxide 
itself  to  the  water,  it  is  much  more  economical,  as  well  as  more  efficient,  to 
produce  it  by  precipitation.  This  is  done  by  simply  adding  to  the  water  some 
aluminium  salt  which  readily  hydrolyzes  (the  sulfate  is  generally  used)  : 

A12(SO4)3  +  6  H2O  =  2  A1(OH)8  +  3  H2SO4 
The  sulfuric  acid  liberated  reacts  with  the  mineral  matter  present  in  the  water, 

Aluminium  chloride  (A1C13).  This  salt  is  prepared  by -passing  dry 
chlorine  or  hydrogen  chloride  through  a  heated  tube  containing  alu- 
minium. It  is  a  crystalline  salt  which  sublimes  without  melting  under 
atmospheric  pressure  and  melts  under  increased  pressure  at  193°.  It 
is  strongly  hygroscopic  and  fumes  when  exposed  to  the  air,  owing  to 
the  formation  of  hydrochloric  acid  through  the  action  of  moisture.  The 
hexahydrate  is  formed  by  dissolving  aluminium  hydroxide  in  hydrochlo1 
ric  acid  and  evaporating  to  crystallization.  When  heated  the  hydrate 
is  decomposed,  forming  aluminium  oxide  and  water,  so  that  the  anhy- 
drous salt  cannot  be  prepared  directly  from  the  hydrate.  The  anhydrous 
salt  is  often  used  for  effecting  certain  syntheses  in  organic  chemistry. 


THE  ALUMINIUM  GROUP  445 

Aluminium  sulfide  (A12S3).  The  sulfide  is  prepared  by  heating  a 
mixture  of  aluminium  and  sulfur  to  a  high  temperature.  It  may  be 
obtained  either  as  a  black  mass  or  as  a  yellow,  crystalline  solid.  It  is 
completely  hydrolyzed  upon  the  addition  of  water : 

A12S8  +  6  H20  =  2  A1(OH)3  +  3  H2S 

When  a  soluble  sulfide, 'such  as  sodium  or  ammonium  sulfide,  is  added 
to  a  solution  of  an  aluminium  salt,  aluminium  sulfide  is  probably 
formed,  but  it  immediately  hydrolyzes,  so  that  the  precipitate  obtained 
consists  of  aluminium  hydroxide.  The  complete  reaction  is  repre- 
sented by  the  following  equation : 

3  Na2S  +  2  A1C13  +  6  H2O  =  2  A1(OH)3  +  6  NaCl  +  3  H2S 

Aluminium  sulfate  (A12(SOJ3).  This  compound  is  prepared  commer- 
cially by  the  action  of  sulfuric  acid  upon  either  bauxite  or  kaolin, 
which  is  an  abundant  silicate  of  aluminium.  In  either  case  the  prep- 
aration of  the  salt  is  somewhat  complicated  by  the  presence  of  im- 
purities in  the  minerals.  The  salt  crystallizes  from  water  in  quite  a 
variety  of  hydrates,  the  usual  one  having  the  composition  expressed 
by  the  formula  A12(SO4)3  •  18  H2O.  It  is  the  cheapest  of  the  soluble 
salts  of  aluminium  and  is  therefore  the  one  most  largely  used  when 
a  salt  of  this  metal  is  desired.  Its  principal  uses  are  in  the  manufac- 
ture of  alum  and  paper,  in  the  purification  of  water,  and  as  a  mordant 
in  dyeing. 

The  sizing  of  paper.  All  paper  intended  for  writing  or  printing  must  be 
sized,  that  is,  coated  over  with  some  substance  that  will  prevent  the  ink  from 
spreading.  Different  methods  have  been  devised  for  doing  this,  the  following 
being  a  common  one :  The  paper  pulp  is  mixed  with  a  soap  made  by  heating 
resin  with  sodium  hydroxide.  A  solution  of  aluminium  sulfate  is  then  added. 
This  reacts  with  the  soap  to  form  sodium  sulfate  and  aluminium  resinate,  the 
latter  compound  being  largely  hydrolyzed  to  resin  and  aluminium  hydroxide. 
In  this  way  the  resin  is  thoroughly  incorporated  with  the  pulp.  As  the  pulp  is 
run  over  hot  rolls  in  making  the  paper,  the  resin  is  melted,  and  on  cooling  forms 
a  thin,  impervious  layer  upon  the  surface  of  the  paper. 

The  alums.  If  solutions  of  aluminium  sulfate  and  potassium  sulfate 
are  mixed  together  and  evaporated,  well-formed  octahedral  crystals 
are  deposited.  These  have  the  composition  expressed  by  the  formula 
K2SO4  •  A12(SO4)3  •  24  H2O.  This  compound  is  a  typical  member  of  a 
class  of  compounds  known  as  the  alums.  The  composition  of  an  alum 
is  expressed  by  the  general  formula  M'2SO4  •  M'"2(SO4)3  •  24  H2O  or, 
more  simply,  as  M'M"'(SO4)2  •  12  H2O,  in  which  M'  represents  a 


446  GENERAL  CHEMISTRY 

univalent  metal  and  Mm  a  trivalent  metal.  For  the  univalent  metal 
one  may  have  any  of  the  alkali  metals  excepting  lithium ;  also  ammo- 
nium, silver,  or  thallium.  For  the  trivalent  metals,  in  addition  to 
aluminium,  one  may  have  iron  or  chromium,  as  well  as  a  few  of  the 
other  metals  which  can  be  obtained  in  a  trivalent  condition.  Those 
alums  that  contain  aluminium  are  white  solids,  those  containing  chro- 
mium have  a  ruby  red  or  purple  color,  while  those  containing  iron 
have  a  violet  tint.  They  all  crystallize  readily,  forming  beautiful 
octahedral  crystals  which  are  isomorphous.  When  heated,  the  water 
of  crystallization  is  evolved,  generally  with  some  sulfur  trioxide, 
leaving  a  residue  known  as  burnt  alum. 

Potassium  alum  (KA1(SO4)2  •  12  H2O)  and  ammonium  alum 
(NH4A1(SO4)2-12H0O)  are  the  most  widely  used.  The  former  has 
the  more  extensive  use.  It  is  sometimes  prepared  from  the  natural 
mineral  alunite  (K2SO4  •  A12(SO4)3  •  4  A1(OH)3).  More  often  it  is 
prepared  by  combining  aluminium  sulfate,  obtained  from  bauxite  or 
kaolin,  with  potassium  sulfate  from  the  Stassfurt  deposits. 

Since  the  alums  crystallize  so  readily,  it  is  easy  to  obtain  them  in  a 
pure  condition.  For  this  reason  the  aluminium  alums  have  long  been 
used  in  place  of  the  much  cheaper  aluminium  sulfate  which  is  difficult 
to  purify.  Improved  methods  have  been  devised  for  preparing  this 
latter  compound,  so  that  now  it  is  taking  the  place  of  alum  for  many 
commercial  uses,  such  as  the  purification  of  water.  The  alums  are 
used  in  the  manufacture  of  paper,  in  water  purification,  and  as  mor- 
dants in  dyeing.  Smaller  quantities  are  used  in  baking  powders  and 
in  certain  foods,  such  as  pickles,  since  it  makes  them  more  crisp. 

Aluminium  carbonate.  The  normal  carbonate  of  aluminium  has  not 
been  prepared.  One  would  naturally  not  expect  it  to  be  formed  in  the 
presence  of  water,  for,  being  a  salt  of  a  very  weak  acid,  as  well  as  of  a 
weak,  insoluble  base,  it  would  be  completely  hydrolyzed.  Accordingly, 
when  a  solution  of  a  carbonate,  such  as  sodium  carbonate,  is  added 
to  a  solution  of  an  aluminium  salt,  carbon  dioxide  is  evolved  and 
aluminium  hydroxide,  mixed  with  small  percentages  of  basic  carbon- 
ates, is  precipitated.  The  main  reaction  takes  place  according  to  the 
following  equation: 

3  Na2CO8  +  2  A1C18  +  3  H2O  =  2  Al(OH)g  +  6  NaCl  +  3  CO2 

Because  of  this  property  alum  is  often  used  as  a  constituent  of  some 
varieties  of  baking  powders. 


THE  ALUMINIUM  GROUP  447 

Alum  baking  powders.  These  consist  of  a  mixture  of  sodium  bicarbonate, 
starch,  and  either  an  alum  or  a  calcined  mixture  of  sodium  and  aluminium  sul- 
fates.  Formerly  both  the  potassium  and  the  ammonium  alum  were  used,  but 
they  have  been  almost  entirely  superseded  by  the  mixture 'just  mentioned,  as 
this  is  cheaper.  This  mixture  is  sold  on  the  market  as  cream  of  tartar  substitute  or 
simply  as  C.T.S.  Upon  the  addition  of  water  to  the  baking  powder  a  reaction 
slowly  takes  place,  resulting  in  the  formation  of  aluminium  hydroxide  and  car- 
bon dioxide,  the  latter  serving  as  the  aerating  agent.  The  starch  absorbs  moisture 
and  thus,  by  preventing  any  reaction,  causes  the  powder  to  retain  its  strength 
until  used.  The  complete  reaction  taking  place  when  water  is  added  to  a  baking 
powder  containing  potassium  alum  is  represented  in  the  following  equation : 

2  KA1(S04)2  +  6  NaHC03  =  2  Al(OH),  +  3  Na2SO4  +  K2SO4  +  6  CO2 

Aluminium  carbide  (A14C3).  This  compound  was  obtained  by  Moissan  by  heating 
aluminium  oxide  with  carbon  in  an  electric  furnace.  Its  chief  interest  lies  in  the 
fact  that  it  reacts  with  water  to  form  methane : 

A14C3  +  12  H20  -  4  A1(OH)3  +  3  CH4 

Since  methane  constitutes  over  90  per  cent  of  natural  gas,  it  has  been  suggested 
that  the  latter  product  may,  in  some  cases  at  least,  have  resulted  from  the  action 
of  water  upon  the  carbides  formed  when  the  crust  of  the  earth  was  in  a  molten 
condition. 

Aluminium  silicates.  In  discussing  the  occurrence  of  aluminium  it 
was  stated  that  the  silicates  of  this  metal  are  widely  and  abundantly 
distributed.  Sometime  in  the  history  of  the  earth's  formation  its  sur- 
face must  have  been  composed  of  a  solid  igneous  rock  formed  by  the 
cooling  of  the  molten  mass.  The  various  silicates  of  aluminium  con- 
stitute by  far  the  largest  percentage  of  these  igneous  rocks.  The  most 
important  of  these  are  the  feldspars,  known  as  orthoclase"(KAlSi3Og), 
albite  (NaAlSi3Og),  and  microcline,  which  has  the  same  chemical  com- 
position as  orthoclase  but  is  different  in  crystalline  structure.  The 
gradual  disintegration,  or  weathering,  of  these  rocks  through  various 
agencies,  such  as  the  action  of  air  and  water,  has  resulted  in  the  for- 
mation of  the  mineral  constituents  of  the  soil.  The  changes  taking 
place  in  this  process  are  often  very  complex  and  not  well  understood. 
Thus,  in  the  weathering  of  orthoclase  the  potassium,  together  with  a 
portion  of  the  silica,  is  removed,  while  at  the  same  time  water  enters 
into  chemical  combination  with  the  residue.  In  this  way  there  is  formed 
the  soft,  plastic  mineral  known  as  kaolinite  (Al0Si2O7  •  2  H2O  or,  as 
often  written,  A12O3  •  2  SiO2  •  2  H2O).  Large  quantities  of  this  mineral 
are  sometimes  found  deposited  in  beds  in  fairly  pure  form.  More  often 
it  has  been  carried  away  by  running  water  and  mixed  with  various 
other  produc.ts  resulting  from  the  crushing  and  weathering  of  rocks, 


448  GENERAL  CHEMISTRY 

especially  silica  (sand)  and  compounds  of  iron,  calcium,  and  magne- 
sium, in  this  way  forming  the  product  known  as  day.  It  is  evident, 
therefore,  that  clay  is  extremely  variable  in  composition,  though  the 
essential  constituent  appears  to  be  kaolinite. 

Ultramarine.  The  mineral  known  as  lapis  lazuli  has  long  been 
known  and  highly  prized  because  of  its  beautiful  blue  color.  The 
powdered  mineral  was  used  by  the  ancients  as  a  color  pigment,  called 
ultramarine.  This  term  is  now  applied  to  the  artificial  product  pre- 
pared by  heating  together,  under  suitable  conditions,  kaolinite,  char- 
coal, sodium  carbonate,  and  sulfur.  The  product  so  obtained  is  very 
similar  in  composition  and  properties  to  the  natural  product  and  is 
much  less  expensive.  By  suitable  variation  of  the  method  of  prepa- 
ration and  of  the  ingredients  employed,  quite  a  variety  of  tints  may 
be  obtained.  It  is  a  very  complex  substance  and  its  exact  composi- 
tion is  not  known  with  certainty.  Large  quantities  of  it  are  used  as 
a  pigment,  especially  in  wall  papers. 

DOUBLE  AND  COMPLEX  SALTS 

In  the  preceding  pages  a  number  of  compounds  have  been  repre- 
sented as  having  formulas  which  indicate  that  they  are  made  up  of  a 
combination  of  two  different  salts.  Most  of  the  Stassfurt  minerals 
(p.  405)  are  of  this  kind ;  the  formula  of  cryolite  is  often  written 
3  NaF  •  AlFg ;  the  fluosilicates  were  formerly  given  such  formulas  as 
2  KF  •  SiF4 ;  finally,  the  alums  present  another  group  of  the  same  gen- 
eral character.  Among  these  salts  two  extreme  cases  can  be  defined. 

Double  salts.  Carnallite  (KC1  •  MgCl2  •  6  H2O)  is  a  good  example 
of  a  double  salt.  When  it  is  dissolved  in  water,  the  solution  acts  as 
though  it  contained  a  mixture  of  potassium  and  magnesium  chlorides, 
both  of  which  are  freely  ionized.  Silver  nitrate  precipitates  all  the 
chlorine  as  silver  chloride  ;  ammonium  phosphate  precipitates  the  mag- 
nesium as  magnesium  ammonium  phosphate  (MgNH4PO4).  A  salt  of 
this  kind,  which  in  solution  decomposes  into  its  constituent  salts  and 
gives  reactions  for  their  individual  ions,  is  called  a  double  salt.  Most 
of  the  Stassfurt  minerals  belong  to  this  class. 

Complex  salts.  Potassium  fluosilicate  (K2SiF6)  may  be  taken  as  an 
example  of  a  complex  salt.  It  can  be  prepared  by  bringing  together 
in  solution  the  fluorides  of  potassium  and  silicon : 


THE  ALUMINIUM  GROUP  449 

The  product  of  the  reaction  cannot  properly  be  regarded  as  a  double 
salt,  for  it  gives  none  of  the  reactions  characteristic  of  the  fluorides, 
from  which  it  is  formed.  Thus,  calcium  chloride,  when  treated  with 
potassium  fluoride,  precipitates  insoluble  calcium  fluoride,  while  with 
fluosilicates  it  enters  into  double  decomposition,  as  shown  in  the 
equation  Q^  +  K2SiF6  =  CaSiF6  +  2  KC1 

The  complex  ion  SiFg~~  acts  as  a  radical  in  all  of  the  reactions  of 
the  fluosilicates.  This  fact,  together  with  the  stability  of  the  salts,  has 
led  chemists  to  regard  the  fluosilicates  as  salts  of  fluosilicic  acid 
(H2SiF6)  rather  than  as  made  up  of  two  different  fluorides  combined, 
as  in  the  case  of  double  salts.  Compounds  of  this  kind,  which  yield 
ions  other  than  those  of  the  salts  from  which  they  may  be  formed,  are 
called  complex  salts. 

Intermediate  types.  The  two  classes  just  described  are  extremes, 
and  there  is  every  gradation  between  them,  the  alums  being  an  exam- 
ple of  this  intermediate  class.  In  rather  concentrated  solution  they 
in  part  dissociate  into  th£  constituent  sulfates,  which  then  give  their 
individual  ions.  In  part  they  act  as  complexes,  giving  the  alkali  metal 
as  one  ion  and  the  remainder  of  the  alum  as  the  other.  These  transi- 
tions constitute  an  equilibrium  which  may  be  represented  thus : 

KA1(S04)2  :<=>:  K+  +  A1(SO4)2~  +=±  Al+  +  +  +  2  SO4~  - 
As  the   solution  becomes  more   and  more   dilute  the    complex   ion 
A1(SO4)2~  tends  to  dissociate  more  completely  into  its  constituent 
ions,  which  are  those  of  a  simple  salt. 

Mixed  salts  and  double  salts.  At  first  sight  the  formula  for  alum  suggests  that 
the  compound  should  be  classified  as  a  mixed  salt  (p.  156)  rather  than  as  a 
double  salt.  The  term  mixed  salt  is,  however,  usually  applied  only  to  those  salts 
in  which  two  different  metals  replace  the  hydrogen  in  one  molecule  of  an  acid, 
as  in  the  case  of  the  salt  NH4MgPO4.  The  formula  for  alum  is  frequently  written 
K2SO4  •  A12(SO4)3  •  24  H2O,  which  indicates  more  clearly  its  characteristic  as  a 
double  salt,  but  does  not  suggest  that  to  some  extent  it  also  acts  as  a  complex  salt. 

GALLIUM,  INDIUM,  AND  THALLIUM 

The  other  members  of  the  aluminium  family  —  gallium,  indium,  and 
thallium  —  are  elements  of  rare  occurrence  in  nature,  and  were  discov- 
ered by  spectroscopic  study  of  various  minerals.  They  have  brilliant 
and  characteristic  spectrum  lines,  and  through  these  are  known  to  be 
widely  distributed  in  certain  classes  of  minerals ;  with  the  exception  of 
a  single  rare  ore  of  thallium  however,  they  have  never  been  found  to 


450  GENERAL  CHEMISTRY 

an  extent  of  more  than  about  0.1  per  cent  in  any  mineral.  All  three 
metals  are  easily  reduced  from  their  oxides,  resembling  zinc  and  lead 
in  this  respect.  Like  aluminium,  they  are  trivalent  in  their  best-known 
compounds,  but  each  of  them  forms  at  least  one  series  of  salts  in 
which  it  is  univalent  or  divalent  as  well.  As  trivalent  metals  each 
forms  a  number  of  alums. 

Gallium.  Gallium  was  discovered  in  certain  zinc  blendes  by  the 
Frenchman  Lecoq  de  Boisbaudran  in  1875,  and  named  in  honor  of 
his  country,  the  Latin  name  for  which  is  G-allia.  It  is  found  in  ores 
of  zinc,  aluminium,  and  iron,  the  richest  known  source  being  the 
iron  from  the  Cleveland  district  in  England,  which  contains  about 
0.003  per  cent  of  gallium.  It  is  a  shining  white  metal  of  very  low 
melting  point  (30.2°).  In  addition  to  the  usual  series  of  salts,  in 
which  it  is  trivalent,  it  forms  a  second  series  in  which  it  is  divalent. 

Indium.  This  metal  was  discovered  in  a  specimen  of  Freiburg  zinc 
blende  by  Reich  and  Richter  in  1863,  while  they  were  examining  it  spec- 
troscopically  for  thallium.  The  name  was  suggested  by  its  character- 
istic indigo-blue  spectrum  line.  Its  richest  ores  do  not  seem  to  contain 
above  0.1  per  cent  indium,  and  it  is  never  found  in  appreciable  quan- 
tities save  in  minerals  containing  zinc.  It  is  a  white  metal,  a  little 
grayer  than  silver,  and  is  as  soft  as  wax.  It  melts  at  155°  and  its 
density  is  about  7.12.  In  addition  to  salts  resembling  those  of  alu- 
minium it  forms  two  other  series,  in  one  of  which  it  is  divalent  and 
in  the  other  univalent. 

Thallium.  Thallium  is  by  far  the  most  abundant  of  these  three  ele- 
ments. It  was  discovered  by  Crookes  in  1861  in  the  slimes  from  the 
lead  chamber  of  a  sulfuric  acid  factory.  The  name  was  suggested  by 
its  brilliant  green  spectrum  line  (thallium  being  derived  from  a  Greek 
word  meaning  "  a  green  twig  ").  It  frequently  accompanies  the  sulfides 
of  the  heavier  metals,  such  as  copper,  lead,  iron,  and  zinc,  and  is  ob- 
tained from  the  flue  dusts  which  are  formed  in  the  roasting  of  such 
sulfides.  To  some  extent  it  also  accompanies  potassium,  and  has  been 
found  in  carnallite.  It  is  a  heavy  metal  (density  11.9)  and  is  softer  than 
lead,  which  it  very  much  resembles  in  appearance.  It  melts  at  301°. 

In  its  chemical  conduct  it  is  very  interesting.  As  a  trivalent  ele- 
ment its  salts  resemble  those  of  aluminium  in  a  general  way,  though 
its  hydroxide  is  rusty  red  in  color,  like  ferric  hydroxide,  and  its  sul- 
fide  (T12S3)  is  soluble  in  ammonium  sulfide,  like  the  sulfides  of  arsenic 
and  antimony.  As  a  univalent  element  it  forms  a  hydroxide  T1OH, 


THE  ALUMINIUM  GROUP  451 

which  is  a  strong,  soluble  base  like  potassium  hydroxide,  and  its  salts 
are,  as  a  rule,  very  similar  to  those  of  potassium.  Its  chloride  and 
cyanicle,  however,  are  insoluble  in  water  and  acids,  resembling  the 
corresponding  compounds  of  silver  in  this  respect.  Thallium  also 
forms  many  double  and  complex  salts. 

THE  RARE  EARTHS 

History.  In  1794  Gadolin  discovered  a  new  mineral,  now  called 
gadolinite,  in  the  mines  of  Ytterby,  near  Stockholm,  and  found  it  to 
contain  an  oxide  unlike  any  known  at  that  time ;  this  was  named 
yttria.  Within  a  few  years  other  minerals  were  found  which  contained 
the  same  oxide,  and  since  then  a  great  many  additional  minerals  have 
been  added  to  the  list.  Almost  at  once  yttria  was  recognized  as  a 
complex  substance,  and  from  that  time  to  the  present  day  the  labors 
of  a  great  number  of  chemists  have  been  only  partially  successful  in 
determining  its  composition.  In  all,  sixteen  different  elements  are  now 
clearly  recognized  to  be  present  in  yttria,  and  these  are  collectively 
called  the  rare  earth  metals.  A  list  of  these,  together  with  their 
atomic  weights,  is  as  follows: 

Scandium 44.1  Gadolinium 157.3 

Yttrium 89.0  Terbium 159.2 

Lanthanum 139.0  Dysprosium 162.5 

Cerium 140.25  Holmium 163.5 

Praseodymium 140.6  Erbium 167.7 

Neodymium 144.3  Thulium 168.5 

Samarium      .......  150.4  Neoytterbium 172.0 

Europium 152.0  Lutecium 174.0 

Occurrence.  Minerals  containing  these  earths  have  been  found  in 
many  different  countries.  In  the  United  States  they  are  found  in  a  num- 
ber of  different  places.  Traces  of  some  of  these  elements,  for  which 
delicate  spectroscopic  tests  are  known,  are  found  to  be  widely  dis- 
tributed in  many  minerals,  but  in  any  appreciable  quantities  they  are 
only  of  very  local  occurrence.  As  far  as  is  known,  no  one  of  them  ever 
occurs  by  itself,  and  a  mineral  which  contains  one  is  likely  to  contain 
most  of  them.  One  of  these  minerals,  monazite,  occurs  largely  in 
North  Carolina  and  Brazil  and  usually  contains  from  1  to  8  per  cent 
of  thorium  oxide,  being  the  chief  source  of  this  rare  and  valuable 
substance.  In  extracting  thorium  large  quantities  of  the  rare  earths 
accumulate,  and  the  supply  available  for  study  is  now  unlimited. 


452  GENERAL  CHEMISTRY 

General  characteristics.  A  number  of  reactions  are  known  which 
separate  these  elements  as  a  group  more  or  less  completely  from  all 
others,  but  no  one  of  the  rare  earths,  with  the  exception  of  cerium, 
can  be  separated  from  the  others  by  a  single  precipitation  of  the  usual 
kind.  Separations  are  only  effected  by  fractional  recrystallizations  or 
precipitations  repeated  hundreds  and  often  thousands  of  times,  so  that, 
with  few  exceptions,  it  is  not  certain  that  the  compounds  of  any  of 
these  elements  have  been  prepared  in  entirely  .pure  form.  .  ( 

These  elements  are  all  trivalent,  though  cerium  is  often  tetravalent 
as  well.  In  general  they  resemble  aluminium,  but  their  hydroxides  are 
stronger  bases,  then-  salts  are  less  hydrolyzed,  and  they  form  no  alums. 
Most  of  their  salts  are  colorless,  though  those  of  neodymium  are  pink, 
those  of  praseodymium  green,  and  the  eerie  salts  yellow  or  red. 

Relation  to  the  periodic  law.  In  atomic  weight  these  elements,  with 
the  exception  of  scandium  and  yttrium,  form  a  continuous  series 
extending  from  the  weight  139  (La)  to  174  (Lu),  the  first  two  in  the 
list  having  much  smaller  weights.  If  they  were  arranged  in  the  periodic 
table  in  the  same  way  as  other  elements,  they  would  evidently  be  dis- 
tributed among  all  the  different  families.  They  are  the  most  similar 
of  all  the  elements,  however,  and  are  therefore  placed  in  the  position 
of  a  single  element  in  the  third  group,  with  the  frank  admission  that 
this  introduces  into  the  table  an  irregular  and  arbitrary  feature.  A 
further  study  of  their  relations  will  probably  throw  much  light  upon 
the  real  meaning  of  the  periodic  law. 

Application.  A  few  practical  applications  have  been  found  for  some 
of  these  substances.  By  the  electrolysis  of  a  mixture  of  a  number  of 
their  compounds  an  alloy  of  the  metals  is  obtained  known  as  mixed 
metal  (mischmetall).  The  heat  of  combustion  of  this  alloy  is  greater 
than  that  of  aluminium,  and  it  is  a  more  powerful  reducing  agent.  It 
is  therefore  sometimes  used  in  place  of  aluminium  in  the  Goldschmidt 
process.  An  alloy  of  cerium  with  iron,  known  as  Auer  metal,  produces 
brilliant  sparks  when  drawn  across  a  rough  surface,  and  owing  to  this 
property  it  is  used  in  the  manufacture  of  gas-lighters.  A  small  quan- 
tity of  cerium  oxide  is  an  essential  constituent  of  gas  mantles,  and 
some  salts  of  cerium  are  used  in  medicine.  Some  of  the  mixed  oxides, 
together  with  zirconium  oxide,  constitute  the  glow  material  in  the 
Nernst  lamp.  Other  applications  are  being  sought,  since  at  the  present 
time  many  tons  of  this  material  annually  go  to  waste. 


CHAPTER  XXXIII 

THE  SILICATE  INDUSTRIES 

The  ceramic  industries.  There  are  a  considerable  number  of  in- 
dustries which  are  based  upon  the  use  of  clay,  sand,  limestone,  and 
feldspar  in  varying  degrees  of  purity,  and  in  as  far  as  they  involve 
chemical  transformations  they  are  closely  allied.  To  a  greater  or 
less  extent  they  depend  upon  the  formation  of  silicates  from  the 
materials  named,  the  bases  being  chiefly  oxides  of  sodium,  potassium, 
calcium,  magnesium,  aluminium,  and  'iron.  These  industries  are 
often  designated  collectively  as  the  ceramic  industries.  They  may  be 
roughly  grouped  into  three  classes,  according  to  whether  they  are 
most  intimately  related  to  the  manufacture  of  glass,  cement,  or  clay 
products. 

Glass.  A  glass  is  essentially  a  material  which,  on  cooling  from  the 
state  of  a  viscous  liquid,  has  failed  to  crystallize  and  yet  has  become 
a  rigid  body.  Pure  quartz,  when  fused  and  cooled,  is  an  example 
of  the  simplest  of  glasses.  The  ordinary  commercial  varieties  of  glass 
are  mixtures  of  various  silicates,  together  with  excess  of  silica.  When 
melted  these  all  mix  together  to  form  a  homogeneous  liquid,  and 
when  this  is  cooled,  it  gradually  hardens  to  a  glass. 

1.  Ingredients  of  glass.  The  ingredients  ordinarily  employed  in  glass- 
making  are  sand,  limestone,  and  the  carbonate  or  sulfate  of  sodium. 
When  a  mixture  of  these  materials  in  the  proper  proportion  is  heated 
to  fusion,  the  volatile  anhydrides  are  driven  out  by  the  silica  (p.  355) 
and  the  bases  remain  in  the  form  of  silicates.  For  glasses  of  fine 
quality  pure  materials  must  be  used,  while  for  cheap  bottle  glass 
ordinary  sand,  limestone,  and  soda  ash  will  serve.  When  sodium  sul- 
fate is  used  in  place  of  sodium  carbonate,  carbon  is  added  to  reduce 
the  sulfate  to  sulfite,  which  is  more  easily  decomposed  by  the  silica. 
Arsenic  trioxide  is  sometimes  added,  and  may  act  either  as  an  oxidiz- 
ing or  a  reducing  agent,  depending  upon  the  conditions.  It  also  forms 
bubbles  on  volatilization  and  thus,  in  some  cases,  may  assist  in  stirring 
the  liquid  and  in  collecting  the  smaller  bubbles  of  other  gases,  which 
are  always  present  in  the  melt. 

453 


454 


GENERAL  CHEMISTRY 


2.  Varieties  of  glass.  By  the  proper  selection  and  proportioning  of 
the  ingredients  a  great  variety  of  glasses  can  be  made.  Ordinary 
window  glass  is  essentially  a  sodium-calcium-magnesium  glass,  and  as 
a  rule  its  composition  closely  approximates  the  percentages  required 
by  the  formula  Na2O  •  CaO  •  6  SiO2.  Since  it  is  made  from  ordinary 
limestone,  which  always  contains  magnesium,  a  variable  percentage 

of  magnesium  oxide  replaces  a  correspond- 
ing percentage  of  calcium  oxide  in  the 
formula.  In  the  harder  glass,  made  to 
resist  chemical  reagents  (Resistance  and 
Jena  glass)  the  sodium  is  largely  replaced 
by  potassium.  The  addition  of  lead  oxide 
makes  a  soft  glass  but  one  which  is  very 
brilliant  and  has  a  high  refractive  index 
suitable  for  some  optical  purposes.  Alumin- 
ium oxide  makes  the  glass  workable  in  the 
blowpipe.  The  oxides  of  barium,  zinc,  and 
boron  are  added  for  special  purposes. 


FIG.  141 


Molding  and  blowing  of  glass.  The  way  in  which 
the  melted  mixture  is  handled  in  the  glass  factory 
depends  upon  the  character  of  the  article  to  be  made.  Many  articles,  such  as 
bottles,  are  made  by  blowing  the  plastic  glass  into  hollow  molds  of  the  desired 
shape.  The  mold  is  first  opened,  as  shown  in  Fig.  141.  A  lump  of  plastic  glass  A 
on  the  hollow  rod  B  is  lowered  into  the  mold,  which  is  then  closed  by  the  han- 
dles C.  By  blowing  into  the  tube  the  glass  is  expanded  into  the  shape  of  the 
mold.  The  mold  is  then  opened  and  the  bottle  lifted  out.  The  neck  of  the  bottle 
must  be  cut  off  at  the  proper  place  and  the  sharp  edges  rounded  off  in  the  flame. 
Other  objects,  such  as  lamp  chimneys,  are  made  by  getting  a  lump  of  plastic 
glass  on  the  end  of  a  hollow  iron  rod  and  blowing  it  into  the  desired  shape 
without  the  help  of  a  mold,  great 
skill  being  required  in  the  ma- 
nipulation of  the  glass.  Window 
glass  is  made  by  blowing  large 
hollow  cylinders  about  6  ft.  long 
and  1^  ft.  in  diameter.  These  are 
cut  longitudinally  and  are  then 
placed  in  an  oven  and  heated 

until  they  soften,  when  they  are  flattened  out  into  plates  (Fig.  142).  Plate  glass  is 
cast  into  flat  slabs,  which  are  then  ground  and  polished  to  perfectly  plane  surfaces. 

3.  Color  of  glass.  The  color  of  glass  is  usually  due  to  the  presence 
of  colored  metallic  silicates.  For  example,  ferrous  silicate  colors  the 
glass  green,  while  ferric  silicate  colors  it  yellow  or  brown.  The  green 


FIG>  142 


THE   SILICATE  INDUSTRIES 


455 


color  can  be  changed  to  the  less  objectionable  yellowish  tint  by  the 
addition  of  manganese  dioxide,  which  acts  as  an  oxidizing  agent,  con- 
verting the  ferrous  compounds  into  ferric.  Cobalt  compounds  form 
deep  blue  silicates,  and  many  other  metals  impart  characteristic  colors. 
Sometimes  the  metals  themselves  are  added  and  form  a  colloidal  sus- 
pension. Copper  and  gold  are  added  to  glass  to  produce  a  rich,  ruby- 
red  color.  Selenium  also  gives  a  beautiful  red  color  to  glass  and  is 
used  both  to  produce  this  color  and  to  compensate  for  the  green  of 
ferrous  silicate.  Opaque,  or  milky,  glasses  are  made  by  adding  mate- 
rials which  remain  suspended  as  solids  in  the  melt,  or  which  melt 
along  with  the  glass  but  do  not  mix  with  it.  In  the  latter  case  an 
emulsion  is  formed,  and  the  turbid  glass  remains  opaque  on  cooling. 
Fluor  spar,  cryolite,  bone  ash,  and  tin  oxide  are  used  in  this  way. 

Cement.  The  term  cement  as  ordinarily  used  at  present  is  applied 
to  those  mortars  known  more  specifically  as  the  hydraulic  cements, 
which  possess  the  property  of  hardening  in  water  as  well  as  in  air. 
These  cements  are  silicate  bodies,  usually  very  highly  basic  in  character, 
which,  when  ground  fine  and  mixed  with  water,  undergo  complex  reac- 
tions resulting  in  the  formation  of  a  hard,  rocklike  mass.  A  number 
of  different  classes  of  cements  are  known,  the  most  important  of  which 
is  the  so-called  Portland  cement. 

Composition  of  Portland  cement.  The  essential  ingredients  of  Port- 
land cement,  together  with  the  limits  of  each  ingredient,  are  given  by 
Bleininger  as  follows : 


INGREDIENTS 

MINIMUM 
PER  CENT 

MAXIMUM 
PER  CENT 

SiO         .     .     . 

19 

26 

Al  O 

4 

11 

Fe  O 

2 

5 

-*-  C2W3 

CaO   .     .     .    .-.-...*    v    .    .    .     .  ;  . 

58 

67 

MffO  . 

o 

5 

SO3    .     .     .     .     .     .    .    ".     .     .     .'    .     . 

0 

25 

Na2O  +  K  O    •  .     .     .     .     . 

3  0 

Manufacture  of  Portland  cement.  The  materials  most  commonly 
employed  are  limestone  or  marl  and  clay  or  shale.  In  general,  how- 
ever, any  substance  may  be  used  which  furnishes  the  ingredients  listed 
in  the  above  table.  Among  the  substances  so  used  is  blast-furnace 
slag,  which  is  an  impure  calcium-aluminium  silicate.  The  materials  to 


456  GENERAL  CHEMISTEY 

be  used  are  coarsely  ground  and  then  mixed  together  in  the  proper 
proportions  and  finely  pulverized.  The  resulting  mixture  is  run  into 
a  furnace  and  burned  to  a  temperature  just  short  of  fusion,  at  which 
temperature  it  vitrifies,  forming  a  grayish  mass  known  as  clinker. 
The  process  of  silicate  formation  is  not  as  complete  as  in  the  case  of 
glass,  but  definite  compounds  are  formed,  among  them  being  those 
represented  by  the  formulas  3  CaO  •  SiO2,  2  CaO  •  SiO2,  2  CaO  •  A12O3. 
Finally,  the  clinker  is  ground  to  a  fine  powder.  Gypsum  is  often 
added  in  the  process ;  this  acts  as  a  negative  catalyzer,  retarding 
the  hardening,  or  setting,  of  the  cement. 

The  setting  of  cement.  The  reactions  which  take  place  upon  the 
addition  of  water  to  cement,  and  which  result  in  the  formation  of  a 
hard,  rocklike  mass,  are  not  at  all  thoroughly  understood.  The  com- 
plex substances  apparently  undergo  hydrolysis  when  they  come  in 
contact  with  water.  The  resulting  compounds  unite  with  water  to 
form  crystalline  hydrates,  producing  the  hard,  compact  mass.  The 
process  of  setting  takes  place  best  in  air,  but  when  wholly  or  partially 
completed,  the  mass  may  be  placed  under  water,  since  the  compounds 
present  are  all  insoluble. 

Growing  importance  of  cement.  Cement  is  rapidly  coming  into  use  for  a  great 
variety  of  purposes.  It  is  often  used  in  place  of  mortar  in  the  construction  of 
brick  buildings.  Mixed  with  crushed  stone  and  sand  it  forms  concrete,  which  is 
used  in  foundation  work.  It  is  also  used  in  making  artificial  stone,  terra-cotta 
trimmings  for  buildings,  artificial  stone  walks  and  floors,  and  the  like.  It  is  being 
used  more  and  more  for  making  many  articles  which  were  formerly  made  of  wood 
or  stone,  and  the  entire  walls  of  buildings  are  sometimes  made  of  cement  blocks 
or  of  concrete. 

Clay  products.  The  crudest  forms  of  clay  products,  such  as  porous 
brick  and  draintile,  have  little  chemistry  involved  in  their  manufac- 
ture. Naturally  occurring  clay  is  molded  into  the  required  form,  dried, 
and  then  burned  in  a  kiln,  but  not  to  a  temperature  at  which  the 
materials  soften.  In  this  process  the  nearly  colorless  ferrous  com- 
pounds in  the  clay  are  converted  into  ferric  compounds,  which  give 
the  usual  red  color  to  these  articles.  In  making  vitrified  brick  the 
temperature  is  raised  to  the  point  at  which  fusion  begins,  so  that  the 
brick  is  partially  changed  to  a  kind  of  glass. 

White  pottery.  This  term  is  applied  to  a  variety  of  articles  varying 
from  the  crudest  porcelain  to  the  finest  chinaware.  While  the  processes 
used  in  the  manufacture  of  the  articles  differ  in  details,  fundamentally 


THE   SILICATE  INDUSTRIES  457 

they  are  the  same  and  may  be  described  under  three  heads :  namely, 
(1)  the  preparation  of  the  body  of  the  ware,  (2)  the  process  of  glazing, 
and  (3)  the  decoration. 

1.  The  body  of  the  ware.   The  materials  used  consist  of  an  artificially 
compounded  clay  made  from  kaolin,  plastic  clay,  and  pulverized  feld- 
spar.   This  mixture  is  plastic  and  is  worked  into  the  desired  shape  by 
molds  or  on  a  potter's  wheel.    The  ware  is  then  dried  and  burned 
until  vitrified,  and  in  this  form  is  known  as  Usque.    This  is  usually 
porous  and  hence  must  be  glazed  to  render  it  nonabsorbent. 

2.  The  glaze.  The  glaze  is  a  fusible  glass  which  is  melted  over  the 
surface  of  the  body.   The  constituents  of  the  glass  are  quartz,  feldspar, 
and  various  metallic  oxides,  often  mixed  with  a  little  boric  oxide. 
These  materials  are  finely  ground  and  mixed  with  water  to  a  paste. 
Sometimes  they  are  first  fused  into  a  glass,  which  is  then  powdered 
and  made  into  the  paste.    The  bisque  is  dipped  into  the  glaze  paste, 
dried,  and  fired  until  the  glaze  materials  melt  and  flow  evenly  over 
the  surface.    The  glaze  must  be  so  chosen  as  to  resist  the  reagents  to 
which  it  is  to  be  exposed,  and  it  must  have  the  same  coefficient  of 
expansion  as  the  body;  otherwise  it  will  check  or  crackle  when  the 
vessel  is  exposed  to  changes  of  temperature.    The  calculation  of  a 
glaze  for  a  given  body  evidently  requires  a  very  thorough  knowledge 
of  the  physical  constants  of  the  clay  from  which  the  body  was  made,  as 
well  as  of  the  properties  contributed  to  the  glaze  by  each  ingredient. 

3.  The  decoration.  If  the  article  is  to  be  decorated,  the  design  may 
either  be  painted  upon  the  body  before  glazing,  when  it  is  said  to  be 
underglazed,  or  it  may  be  painted  upon  the  glaze  and  the  article  fired 
again,  the  pigments  melting  into  the  glaze.    In  the  latter  case  it  is 
said  to  be  overglazed.    In  the  former  case  the  pigments  used  are,  as 
a  rule,  metallic  oxides,  while  in  the  latter  case  they  are  often  colored 
glasses. 


CHAPTER  XXXIV 

THE  IRON  FAMILY 


SYMBOLS 

ATOMIC 
WEIGHT 

DENSITY 

APPROXIMATE 
MELTING  POINT 

OXIDES 

Iron      .     .     . 

Fe 

55.84 

7.86 

1505° 

FeO,  Fe2O3 

Cobalt  .     .     . 

Co 

58.97 

8.6 

1490° 

CoO,  Co2O3 

Nickel  .     .     . 

Ni 

58.68 

8.9 

1452° 

NiO,  Ni2O3 

The  family.  The  elements  iron,  cobalt,  and  nickel  bear  a  relation 
to  one  another  which  is  different  from  that  existing  among  the  mem- 
bers of  any  other  family  as  yet  considered.  Their  atomic  weights  are 
very  close  together,  and  in  the  periodic  table  they  are  placed  in  one 
family,  not  because  the  plan  of  arrangement  brings  them  together, 
but  because  they  are  so  similar  (p.  235)  and  evidently  constitute  a 
natural  family. 

Iron  is  an  abundant  element  (p.  14),  while  the  other  two  are  rather 
rare  and  local  in  their  occurrence.  To  a  limited  extent  they  have  all 
been  found  in  nature  in  the  native  state,  particularly  in  meteorites. 
The  members  of  one  class  of  these  meteorites,  known  as  the  "  meteoric 
irons,"  consist  principally  of  uncombined  iron  mixed  with  smaller  per- 
centages of  free  nickel,  traces  of  cobalt  also  being  present  in  some 
cases.  In  the  combined  state  iron  is  found  in  large  deposits  both  as 
oxides  and  as  sulficles,  while  nickel  and  cobalt  are  found  together 
combined  with  sulfur  or  arsenic  and  usually  associated  with  silver 
and  copper.  Their  salts  are  nearly  all  highly  colored.  Iron  is  note- 
worthy as  the  first  metal  to  be  described  in  detail  which  exerts  two 
different  valences  and  forms  two  series  of  salts.  In  the  one  (the  fer- 
rous) the  iron  is  divalent,  and  in  many  respects  ferrous  salts  resemble 
those  of  magnesium ;  in  the  other  (the  ferric)  it  is  trivalent,  the  ferric 
salts  resembling  those  of  aluminium.  It  is  therefore  appropriate  to 
consider  iron  at  the  present  time.  It  is  also  worthy  of  note  that  the 
metals  which  remain  to  be  considered  are  isolated  by  reduction  of 
their  oxides  or  from  their  sulfides,  and  not  by  electrolytic  methods. 

458 


THE  IRON  FAMILY  459 

IRON 

Occurrence.  Iron  has  long  been  known,  since  its  ores  are  very 
abundant  and  it  is  not  difficult  to  prepare  the  metal  from  them  in 
fairly  pure  form.  It  occurs  in  large  deposits  as  oxides,  sulfides,  and 
carbonates,  and  in  smaller  quantities  in  a  great  variety  of  minerals. 
Indeed,  very  few  rocks  and  soils  are  free  from  small  amounts  of  iron, 
and  it  is  assimilated  by  both  plants  and  animals.  It  is  a  constituent 
of  both  chlorophyll  and  hsemoglobin,  and  plays  a  fundamental  part  in 
life  processes.  The  most  important  ores  are  the  following : 

Hematite  ...     ..    Fe2Os  Siderite  ....     FeCO3 

Magnetite      .     .     .     Fe3O4  Limonite     .     .     .     2  Fe2O3  •  3  H2O 

While  iron  ore  is  mined  in  a  number  of  different  localities  in  the 
United  States,  the  great  center  of  production  is  in  the  territory  border- 
ing on  Lake  Superior,  the  ore  being  chiefly  hematite.  Large  amounts 
are  also  mined  near  Birmingham,  Alabama. 

Preparation  of  pure  iron.  Pure  iron  may  be  prepared  in  the  form  of 
a  fine  powder  by  heating  the  oxide  in  a  current  of  hydrogen,  though 
the  product  contains  occluded  hydrogen  unless  the  process  is  carried 
out  at  a  high  temperature.  It  may  be  obtained  in  coherent  masses  by 
the  electrolysis  of  ferrous  sulf ate  between  iron  electrodes.  To  prevent 
the  occlusion  of  hydrogen,  which  makes  the  metal  hard  and  brittle,  it 
is  necessary  to  conduct  the  electrolysis  at  about  100°  and  to  add  some 
calcium  chloride  to  the  electrolyte.  By  such  methods  Burgess  has 
obtained  iron  said  to  be  99.98  per  cent  pure. 

Properties  of  pure  iron.  Pure  iron  is  a  silvery  metal  having  a  density 
of  7.86  and  a  melting  point  of  1505°.  It  is  ductile  and  malleable  and 
is  almost  as  soft  as  aluminium.  Xt  is  especially  well  adapted  to  the 
manufacture  of  electromagnets,  since  it  acquires  and  loses  magnetic 
properties  much  more  rapidly  than  do  the  ordinary  varieties  of  iron. 
It  is  not  acted  upon  by  dry  air  but  rusts  in  moist  air. 

The  iron  of  commerce.  Iron  differs  from  most  of  the  other  metals 
used  in  the  industries  in  that  the  pure  metal  is  rarely  obtained  and  is 
of  limited  value,  while  that  which  contains  small  percentages  of  other 
elements  exhibits  a  wide  variety  of  properties  and  is  of  the  greatest 
importance.  Carbon  is  always  present  in  quantities  which  range  from 
mere  traces  up  to  7  per  cent.  According  to  the  conditions  under  which 
the  metal  is  produced,  this  carbon  may  be  in  the  form  of  graphite  scat- 
tered through  the  iron,  or  as  a  solid  solution  of  carbon  in  iron,  or  in 


460 


GENERAL  CHEMISTRY 


combination  with  the  iron  in  the  form  of  a  carbide.  The  most  im- 
portant of  these  carbides  has  the  formula  Fe3C  and  is  a  hard,  brittle 
substance  known  as  cementite.  Manganese  and  silicon,  together  with 
traces  of  phosphorus  and  sulfur,  are  also  present. 

The  properties  of  iron  are  much  modified  by  the  percentages  of 
these  constituents,,  by  their  form  of  combination  in  the  iron,  and  by 
the  treatment  of  the  metal  during  its  production  from  the  ore.  Owing 
to  these  facts  many  varieties  of  iron  are  recognized  in  commerce,  the 
chief  of  which  are  cast  iron,  wrought  iron,  and  steel. 

Cast  iron.  Ordinarily  the  first  step  in  the  manufacture  of  any  variety 
of  commercial  iron  is  the  production  of  cast  iron.  The  ores,  with  the 

exception  of  the  oxides,  are  first  roasted. 
They  are  then  mixed  with  a  suitable  flux 
and  reduced  by  heating  with  carbon,  usu- 
ally in  the  form  of  coke.  As  a  rule  the  ore 
carries  with  it  minerals  rich  in  silicon  and 
aluminium,  and  in  such  cases  limestone 
is  used  as  a  flux,  the  resulting  slag  being 
essentially  a  calcium-aluminium  glass. 

Blast-furnace  process.  The  reduction  is  car- 
ried out  in  a  large  tower  called  a  blast  furnace 
(Fig.  143).  This  is  usually  80  ft.  high  and  20  ft. 
in  internal  diameter  at  its  widest  part,  narrow- 
ing somewhat  both  toward  the  top  and  bottom. 
The  walls  are  built  of  steel  and  are  lined  with 
fire  brick.  The  base  is  provided  with  a  number 
of  pipes  A  called  tuyeres,  through  which  hot 
air  is  forced  into  the  furnace.  The  tuyeres  are 
supplied  from  a  large  pipe  B,  which  circles  the 
furnace  as  a  girdle.  At  the  base  of  the  furnace 
is  an  opening  through  which'  the  liquid  metal 
can  be  drawn  off  from  time  to  time.  There  is 
also  a  second  opening  C,  somewhat  above  the 
—  first,  through  which  the  excess  of  slag  over- 
flows. The  top  is  closed  by  a  movable  trap  /), 
called  the  cone,  and  through  this  the  materials 

to  be  used  are  introduced.  The  gases  resulting  from  the  combustion  of  the  fuel 
and  the  reduction  of  the  ore,  together  with  the  nitrogen  of  the  air  admitted 
through  the  tuyeres,  escape  through  pipes  E.  These  gases  are  very  hot  and  con- 
tain enough  carbon  monoxide  to  be  combustible ;  they  are  accordingly  utilized 
for  heating  the  blast  admitted  through  the  tuyeres,  and  as  fuel  for  the  engines. 
Charges  consisting  of  coke,  ore,  and  flux  in  proper  proportion  are  introduced 
into  the  furnace  at  intervals  through  the  cone.  The  coke  burns  fiercely  in  the 


FIG.  143 


THE  IKON  FAMILY  461 

hot-air  blast,  forming  carbon  dioxide,  which  is  at  once  reduced  to  carbon  monoxide 
as  it  passes  over  the  highly  heated  carbon.  The  temperature  of  the  furnace  at 
the  point  at  which  the  hot  air  enters  is  about  1600°,  but  gradually  decreases 
toward  the  top  of  the  furnace,  at  which  it  is  only  from  300°  to  400°.  Reduction 
of  the  ore  begins  at  the  top  of  the  furnace  through  the  action  of  the  carbon 
monoxide.  As  the  ore  slowly  descends  the  reduction  is  completed  and  the  result- 
ing iron  melts  and  collects  as  a  liquid  in  the  bottom  of  the  furnace,  the  lighter 
slag  floating  above  it.  After  a  considerable  quantity  of  iron  has  collected,  the 
slag  is  drawn  off  through  C,  and  the  iron  is  run  out  into  ladles  and  taken  to  the 
converters  for  the  manufacture  of  steel;  or  it  is  run  into  sand  molds  and  cast 
into  ingots  called  pigs.  The  process  is  a  continuous  one,  and  when  the  furnace 
is  once  started,  it  is  kept  in  operation  for  months  or  even  years  without  inter- 
ruption. The  iron  is  withdrawn  at  intervals  of  about  six  hours. 

Properties  of  cast  iron.  The  product  of  the  blast  furnace  is  called 
cast  iron.  It  varies  considerably  in  composition,  but  always  contains 
over  2  per  cent  of  carbon,  variable  amounts  of  silicon,  and  at  least 
traces  of  phosphorus  and  sulfur.  Two  extreme  varieties  of  cast  iron 
are  recognized :  namely,  gray  iron  and  white  iron.  In  gray  iron  the 
carbon  is  present  partly  in  the  form  of  cementite  and  partly  as 
graphite,  the  latter  of  which  gives  the  metal  its  gray  color.  In  white 
cast  iron  almost  all  the  carbon  is  in  the  combined  state  in  the  form 
of  cementite ;  hence,  this  variety  is  much  harder  and  more  brittle 
than  the  gray  iron.  Between  these  two  extreme  types  there  are  all 
intermediate  varieties.  Cast  iron  is  hard  and  brittle  and  melts  at 
about  1100°.  It  cannot  be  welded  or  forged,  but  is  easily  cast  in  sand 
molds.  It  is  rigid  but  not  elastic,  and  its  tensile  strength  is  small. 
It  is  used  for  making  castings  and  in  the  manufacture  of  other 
varieties  of  iron. 

Wrought  iron.  Wrought  iron  is  made  from  cast  iron  by  burning  out 
most  of  the  carbon,  silicon,  phosphorus,  and  sulfur  which  it  contains. 

The  process  is  carried  out  in  a  puddling  furnace.  The  floor  of  the  furnace  is 
covered  with  a  layer  of  iron  oxide,  and  on  this  is  placed  the  charge  of  cast  iron, 
together  with  some  suitable  flux  (generally  limestone).  The  fuel  is  burned  in  a 
fire  box  at  the  side  of  the  furnace,  and  the  flame  is  led  over  the  charge  of  cast 
iron,  the  heat  being  reflected  down  upon  it  by  a  low,  arching  roof.  The  iron  is 
soon  melted,  and  the  sulfur,  phosphorus,  and  silicon  are  oxidized  by  the  iron 
oxide,  forming  acid  anhydrides,  which  combine  with  the  flux  or  with  the  iron 
oxide  to  form  a  slag.  The  carbon  is  also  oxidized  and  escapes  as  carbon  dioxide. 
As  the  iron  is  freed  from  other  elements  it  becomes  pasty,  owing  to  the  higher 
melting  point  of  the  purer  iron,  and  in  this  condition  forms  small  lumps,  which 
are  raked  together  into  a  larger  one.  The  large  lump  is  then  removed  from  the 
furnace  and  rolled  or  hammered  into  bars,  most  of  the  slag  being  squeezed  out 
in  this  process. 


462  GENERAL  CHEMISTRY 

Properties  of  wrought  iron.  Wrought  iron  has  a  fibrous  structure, 
being  composed  of  fibers  of  pure  iron  (which  is  known  as  ferrite)  sepa- 
rated by  slag.  The  amount  of  slag  present  varies  from  0.1  per  cent 
to  2  per  cent.  The  ferrite  present  contains  less  than  0.3  per  cent  of 
carbon  and  not  more  than  traces  of  other  elements.  Wrought  iron 
is  soft,  malleable,  and  ductile.  While  its  tensile  strength  is  greater 
than  cast  iron,  it  is  less  than  that  of  most  steel.  Its  melting  point  is 
much  higher  than  that  of  cast  iron.  -Wrought  iron  is  no  longer  pro- 
duced to  the  same  relative  extent  as  in  former  years,  since  soft  steel 
can  be  made  at  less  cost  and  has  almost  the  same  properties. 

Steel.  Steel,  like  wrought  iron,  is  made  from  cast  iron  by  burning 
out  a  part  of  the  carbon,  silicon,  phosphorus,  and  sulfur  which  it  con- 
tains, but  the  processes  used  are  different  from  that  employed  in  the 
manufacture  of  wrought  iron.  Nearly  all  the  steel  of  commerce  is 
made  by  one  of  two  general  methods,  known  as  the  Bessemer  process 
and  the  open-hearth  process.  There  are  two  modifications  of  each  of 
these  processes,  based  upon  the  differences  in  the  material  used  in 
the  lining  of  the  furnaces:  (1)  In  the  one  the  furnaces  are  lined 
with  silica,  which  is  an  acid  anhydride.  This  modification  is  therefore 
known  as  the  acid  process.  In  the  steel  made  in  these  furnaces  the 
carbon  and  silicon  are  removed,  but  all  of  the  phosphorus  and  sulfur 
in  the  original  cast  iron  are  retained,  since  no  fluxing  material  adapted 
to  their  removal  is  present.  The  acid  process  is  employed  when  the 
cast  iron  to  be  used  is  very  low  in  phosphorus  and  sulfur.  (2)  In  the 
other  modification  the  furnace  is  lined  with  limestone  or  dolomite,  and 
this  modification  is  known  as  the  basic  process.  In  such  furnaces  both 
the  phosphorus  and  the  sulfur  are  removed,  together  with  the  carbon 
and  silicon.  These  furnaces  are  therefore  employed  when  the  cast  "iron 
contains  appreciable  percentages  of  phosphorus  and  sulfur.  Practi- 
cally all  of  the  steel  produced  in  the  United  States  is  made  by  either 
the  acid  Bessemer  or  the  basic  open-hearth  process.  A  brief  description 
of  these  methods  follows. 

Acid  Bessemer  process.  This  process,  invented  about  1880,  is  carried  out  in 
great  egg-shaped  crucibles  called  converters  (Fig.  144),  each  one  of  which  will 
hold  as  much  as  15  tons  of  steel.  The  converter  is  built  of  steel  and  lined  with 
silica.  It  is  mounted  on  trunnions,  so  that  it  can  be  tipped  over  on  its  side  for 
filling  and  emptying.  One  of  the  trunnions  is  hollow,  and  a  pipe  connects  it 
with  an  air  chamber  A,  which  forms  a  false  bottom  to  the  converter.  The  true 
bottom  is  perforated,  so  that  air  can  be  forced  in  by  an  air  blast  admitted  through 
the  trunnion  and  the  air  chamber. 


THE  IKON  FAMILY 


463 


White-hot  liquid  cast  iron  from  a  blast  furnace  is  run  into  the  converter 
through  its  open,  necklike  top  B,  the  converter  being  tipped  over  to  receive  it ; 
the  air  blast  is  then  turned  on  and  the 
converter  turned  to  a  nearly  vertical 
position.    The    carbon    and  silicon   in 
the  iron  are  rapidly  oxidized  (the  sili- 
con first   and   then   the    carbon),  the 
oxidation  being  attended  by  a  brilliant 
flame.   The  heat  of  the  reaction,  largely 
due  to  the  combustion  of  silicon,  keeps 
the  iron  in  a  molten  condition.    The  air 
blast  is  continued  until  the  character  of 
the  flame  shows  that  all  the  carbon  has  been  burned  away. 
The  process  requires  from  15  to  20  minutes,  and  when  it  is 
complete,  the  desired  quantity  of  carbon  (generally  in  the 

form  of  high   carbon 

iron  alloys)  is  added 

and    allowed   to    mix 

thoroughly  with    the 

fluid.  The  converter  is 

then  tilted  (Fig.  145), 

and  the  steel  run  into 

molds,  and  the  ingots  FIG.  144 

so    formed    are    ham- 
mered or  rolled  into  rails  or  other  objects. 
Basic  open-hearth  process.  Fig.  146  shows 
FIG.  145  the  simpler  parts  of  the  type  of  furnace  used 

in  this  process.    The  hearth  of  the  furnace 

is  about  40  ft.  in  length,  12  ft.  in  width,  and  2  ft.  in  depth,  and  is  lined  with 
limestone  or  dolomite  (A A}.  Either  gas  or  sprayed  oil  is  used  as  fuel.  Below 
the  furnace  is  placed  a  checkerwork  of  brick  so  arranged  that  the  hot  products 


Hot  Air 


FIG.  146 


of  combustion  escaping  from  the  furnace  may  be  conducted  through  it,  thus 
heating  the  bricks  to  a  high  temperature.  Both  the  air  necessary  for  combustion 
and  the  gaseous  fuel  (unless  decomposed  by  heating,  as  in  the  case  of  natural 


464  GENERAL  CHEMISTRY 

gas  and  sprayed  oil)  are  preheated  by  passing  them  over  the  hot  bricks,  so  that 
the  temperature  reached  during  combustion  is  greatly  increased.  The  gas  enter- 
ing through  C  comes  in  contact  at  D  with  the  hot  air  entering  through  B,  and 
a  vigorous  combustion  ensues,  the  flame  passing  above  and  over  the  cast  iron 
and  lime  with  which  the  furnace  is  charged.  The  products  of  combustion  escape 
through  E  and  F.  At  the  temperature  reached  the  carbon  in  the  cast  iron  is 
removed  in  the  form  of  the  oxide,  the  escaping  gas  giving  the  melted  metal  the 
appearance  of  boiling.  The  silicon,  phosphorus,  and  sulfur  unite  with  oxygen  to 
form  acid  anhydrides ;  these  combine  with  the  lime  to  form  a  slag,  and  this  rises 
to  the  surface  of  the  melted  charge  and  is  easily  removed.  When  a  test  shows 
the  desired  percentage  of  carbon  present,  the  melted  steel  is  run  into  large  ladles 
and  then  into  molds.  An  average  furnace  produces  about  50  tons  of  steel  in  a 
given  charge,  approximately  8  hours  being  required  in  the  process.  At  present 
by  far  the  largest  amount  of  steel  produced  in  the  United  States  is  made  by 
this  process. 

Properties  of  steel.  Steel  contains  from  a  trace  up  to  2  per  cent  of 
carbon,  less  than  0.1  per  cent  of  silicon,  and  not  more  than  traces  of 
phosphorus  and  sulfur.  When  desired,  steel  may  be  made  so  pure 
that  it  contains  only  traces  of  carbon.  Indeed,  a  product  containing 
99.85  per  cent  of  iron  is  now  being  made  by  the  open-hearth  process. 
Such  steel  is  very  soft.  As  the  carbon  content  rises  the  steel  becomes 
harder  and  less  ductile  ;  at  the  same  time  the  tensile  strength  increases 
until  the  carbon  amounts  to  about  1.1  per  cent,  after  which  it  decreases. 

Relation  of  the  three  varieties  of  iron.  Wrought  iron  consists  of 
fibers  of  nearly  pure  iron  (ferrite)  separated  by  traces  of  slag,  while 
most  steel  contains  an  appreciable  amount  of  alloy  material  (chiefly 
carbon)  and  cast  iron  contains  still  more  of  the  same  substances.  It  is 
impossible,  however,  to  assign  a  given  sample  of  iron  to  one  of  these 
three  classes  on  the  basis  of  its  chemical  composition  alone.  For  ex- 
ample, a  low  carbon  steel  may  contain  less  carbon  than  a  given  sample 
of  wrought  iron.  The  classification  of  commercial  iron  into  cast  iron, 
wrought  iron,  and  steel  is  really  based  on  the  method  of  manufacture. 
The  product  of  the  blast  furnace  is  cast  iron,  that  of  the  puddling 
furnace  is  wrought  iron,  that  of  the  Bessemer  and  open-hearth 
processes  is  steel. 

The  hardening  and  tempering  of  steel.  When  steel  containing  from 
0.5  to  1.5  per  cent  of  carbon  is  heated  to  a  relatively  high  temperature 
and  then  cooled  suddenly  by  plunging  it  into  cold  water  or  oil,  it 
becomes  very  hard  and  brittle.  When  gradually  reheated  and  then 
allowed  to  cool  slowly,  this  hardened  steel  becomes  softer  and  less 
brittle,  and  this  process  is  known  as  tempering. 


THE  IRON  FAMILY  465 

By  properly  regulating  the  temperature  to  which  the  steel  is  reheated  in  tem- 
pering, it  is  possible  to  obtain  almost  any  condition  of  hardness  demanded  for  a 
given  purpose,  as  for  making  springs  or  cutting  tools.  Steel  assumes  different 
color  tints  at  different  temperatures,  and  by  these  the  experienced  workman  can 
tell  when  the  desired  temperature  has  been  reached.  Lake  gives  the  following 
temperatures  as  suited  to  the  tempering  of  the  tools  specified : 

220°  —  paper  cutters,  wood-engraving  tools  275°  —  axes,  springs 

240°  —  knife  blades,  rock  drills  290°  —  needles,  screw  drivers 

260°  —  hand-plane  cutters  and  cooper's  tools          300°  —  wood  saws. 

The  changes  which  attend  the  hardening  and  tempering  of  steel  are 
very  complex  and  are  just  beginning  to  be  understood.  We  shall 
simply  note  here  that  the  different  treatments  which  the  iron  receives 
in  the  processes  of  hardening  and  tempering  result  in  a  change  in  the 
condition  of  the  carbon  present,  and  hence  in  a  change  in  the  properties 
of  the  product. 

Steel  alloys.  It  has  been  found  that  small  quantities  of  a  number 
of  different  elements,  when  added  to  steel,  very  much  improve  its 
quality  for  certain  purposes.  Among  the  elements  most  used  in  this 
way  are  manganese,  silicon,  nickel,  chromium,  tungsten,  molybdenum, 
vanadium,  and  titanium.  These  elements  may  act  in  two  different 
ways.  Some  of  them,  such  as  titanium  and  vanadium,  act  mainly  as 
purifiers,  their  function  being  to  remove  any  gases  (chiefly  oxygen  and 
nitrogen)  dissolved  in  the  iron.  Others,  such  as  nickel  and  chromium, 
form  alloys  with  the  steel,  the  properties  of  which  vary  according  to  the 
element  present.  Thus,  nickel  renders  the  steel  harder  and  increases 
its  tensile  strength  and  elastic  limit. 

The  approximate  composition  and  the  uses  of  some  of  these  steel 
alloys  are  as  follows : 

3.5%  nickel armor  plate 

3.5%  nickel  and  3.5%  chromium armor  plate  and  projectiles 

4.0%  manganese      .          burglar-proof  safes 

6.0%  chromium  and  from  8  to  24%  tungsten  .     .  high-speed  lathe  tools 

6.0%  chromium  and  10%  molybdenum  ....  high-speed  lathe  tools 

0.1%  titanium car  rails  and  steel  castings 

0.1%  vanadium automobile  parts 

Passive  iron.  Iron  readily  dissolves  in  both  dilute  and  concentrated  nitric  acid, 
but  when  it  is  brought  into  contact  with  fuming  nitric  acid,  that  is,  with  nitric 
acid  containing  nitrogen  dioxide  in  solution,  it  loses  many  of  its  characteristic 
properties  and  is  then  said  to  be  in  a  passive  state.  For  example,  such  iron  is  no 
longer  attacked  by  dilute  nitric  acid,  nor  does  it  precipitate  copper  and  silver  from 
solutions  of  their  salts,  as  does  ordinary  iron.  The  metal  loses  its  passivity  when 


466 


GENERAL  CHEMISTRY 


FIG.  147 


it  is  rubbed,  scratched,  or  given  a  sharp  blow,  or  when  certain  other  metals 
are  brought  into  contact  with  it.  A  number  of  other  metals,  including  cobalt, 
nickel,  and  chromium,  act  in  a  similar  way.  No  entirely  satisfactory  explanation 
has  been  offered,  to  account  for  this  phenomenon.  According  to  one  assump- 
tion the  nitric  acid  forms  a  thin  protective  film  of  metallic  oxide  over  the  metal. 

This  view  receives  support  from 
the  fact  that  other  oxidizing  agents 
may  be  substituted  for  nitric  acid 
in  rendering  iron  passive.  The  pas- 
sivity of  iron  may  be  illustrated 
in  the  following  simple  way :  A 
piece  of  sheet  iron  A  (Fig.  147) 
is  immersed  for  a  few  moments 
in  fuming  nitric  acid  contained  in  B.  It  is  next  lowered  into  pure  water  in  (7, 
in  order  to  wash  off  the  adhering  acid.  Finally,  it  is  dipped  for  a  moment  into  a 
solution  of  copper  sulfate  in  D.  Apparently  no  change  takes  place.  If  now  the 
iron  is  struck  a  sharp  blow,  it  at  once  loses  its  passivity  and  regains  its  normal 
property  of  replacing  copper  from  copper  sulfate,  as  is  evidenced  by  the  rapid 
formation  of  a  thin  film  of  the  metal  over  the  entire  surface  of  the  iron. 

The  rusting  of  iron.  A  number  of  different  theories  have  been  advanced  to 
account  for  the  changes  taking  place  in  the  rusting  of  iron.  The  most  satisfac- 
tory of  these  is  known  as  the  electrolytic  theory.  According  to  this  the  primary 
reaction  in  the  rusting  of  iron  is  between  iron  and  water,  as  expressed  in  the 
following  equation : 

Fe  +  2  (H+,  OH-) *•  Fe+  +,  2  OH-  +  H2 

The  ions  Fe+  +  and  2  OH~  then  combine  to  form  ferrous  hydroxide  (Fe(OH)2). 
This  is  further  acted  upon  by  oxygen  and  moisture,  and  forms  the  complex  sub- 
stance known  as  iron  rust.  It  is  evident  that  the  composition  of  rust  will  vary 
according  to  the  conditions  of  its  formation. 

Compounds  of  iron.  The  compounds  of  iron  are  much  more  numer- 
ous than  those  of  any  metal  so  far  considered,  for  not  only  does  iron 
form  two  series  of  simple  salts,  but  it  is  a  constituent  of  many  com- 
plex salts  as  well.  It  will  be  possible  to  mention  only  a  few  typical 
individuals  in  each  class. 

Ferrous  compounds.  The  ferrous  salts,  resembling  those  of  mag- 
nesium not  only  in  formula  but  often  in  degree  of  hydration,  are 
usually  nearly  white  when  prepared  by  precipitation,  but  are  colored 
some  shade  of  light  green  or  yellow  when  well  crystallized.  They  are 
not  very  greatly  hydrolyzed  in  solution,  since  ferrous  hydroxide  is 
about  as  strong  a  base  as  the  hydroxide  of  magnesium  or  of  zinc,  but 
they  are  readily  oxidized,  as  will  be  explained  later.  The  soluble 
salts  are  most  easily  prepared  by  dissolving  iron  in  the  appropriate 
acid,  the  insoluble  ones  by  precipitation. 


THE  IRON  FAMILY  467 

Ferrous  hydroxide  (Fe(OH)2)  ;  ferrous  oxide  (FeO).   Ferrous  hydrox- 
ide (Fe(OH)2)  forms  as  a  white,  nearly  insoluble  precipitate  when  a 
solution  of  a  ferrous  salt  is  treated  with  a  soluble  base.    On  exposure 
to  the  air  and  moisture  it  quickly  oxidizes  to  ferric  hydroxide  : 
4  Fe(QH)2  +  2  H2O  +  O2  =  4  Fe(OH)3 

When  heated  out  of  contact  with  air,  it  is  converted  into  ferrous 
oxide  (FeO).  The  latter  compound  is  more  easily  obtained  as  a  black, 
combustible  powder  by  heating  ferrous  oxalate  : 


Ferrous  chloride  (FeCl2).  Anhydrous  ferrous  chloride  is  prepared  by 
strongly  heating  iron  in  a  current  of  hydrogen-  chloride,  the  salt  con- 
densing in  the  colder  portions  of  the  tube  in  white,  pearly  scales.  It 
dissolves  in  water,  with  evolution  of  much  heat,  and  from  this  solution 
crystallizes  as  the  green  tetrahydrate  FeCl2  •  4  H2O.  The  latter  salt 
is  more  easily  obtained  by  dissolving  iron  in  hydrochloric  acid  and 
evaporating  the  solution  out  of  contact  with  the  air. 

Ferrous  sulfate  (FeSOJ.  The  sulfate  is  the  most  familiar  ferrous 
salt  and  has  important  uses.  It  is  easily  prepared  by  dissolving  iron 
in  dilute  sulfuric  acid  and  evaporating  to  crystallization.  It  is  then 
obtained  in  large  monoclinic  crystals  of  the  composition  FeSO4  •  7  H2O, 
known  as  green  vitriol,  or  copperas.  In  the  industries  it  is  obtained  from 
the  liquors  which  result  from  cleaning  sheet  steel  with  sulfuric  acid 
preparatory  to  tinning  or  galvanizing  the  steel.  It  is  also  manufac- 
tured by  the  oxidation  of  the  abundant  mineral  pyrites  (FeS2),  usu- 
ally after  a  careful  partial  roasting  which  converts  the  pyrites  into 
ferrous  sulfide  (FeS)  :  FeS  +  2  O2  =  FeSO4 

It  is  used  as  a  preservative,  as  a  reagent  for  killing  weeds,  in  the  dyeing 
industry,  as  a  substitute  for  aluminium  sulfate  in  water  purification 
(p.  444),  and  in  the  manufacture  of  black  inks. 

Inks.  Most  of  the  common  black  inks  are  made  by  treating  an  infusion  of  nut- 
galls  with  ferrous  sulfate.  The  nut-galls  are  rich  in  tannic  acid,  and  this,  with 
ferrous  sulfate,  gives  a  nearly  black  precipitate.  By  the  addition  of  colloidal 
materials,  such  as  gum  arable  or  dextrin,  the  precipitation  is  greatly  delayed, 
though  the  intensely  black  color  develops  at  once.  Some  preservative  is  usually 
added  to  prevent  the  ink  from  molding. 

The  vitriols.  The  term  vitriol  is  applied  to  the  hydrated  sulfates  of 
a  number  of  divalent  metals.  These  compounds  are  of  two  distinct 
types  :  the  one  group  forms  monoclinic  crystals  which  contain  seven 


468  GENERAL  CHEMISTRY 

molecules  of  water  of  crystallization ;  the  other  forms  triclinic  crystals 
with  five  molecules  of  water.  All  the  salts  in  a  given  series  are 
isomorphous,  and  many  of  the  vitriols  are  dimorphous,  crystallizing 
in  both  forms.  The  sulfates  of  iron,  zinc,  and  magnesium  are  the 
most  familiar  representatives  of  the  monoclinic  vitriols,  while  copper 
sulfate  (CuSO4  •  5  H2O)  is  the  best-known  triclinic  vitriol. 

Ferrous  ammonium  sulfate  (Mohr's  salt)  ((NH4)2S04  •  FeS04  •  6  H20). 
When  ammonium  sulfate  and  ferrous  sulfate  are  brought  together 
in  solution  in  molecular  proportions,  a  double  salt  of  the  formula 
(NH4)2SO4  •  FeSO4  •  6  H2O,  known  as  Mohr's  salt,  separates  on  crys- 
tallization. This  salt  oxidizes  less  readily  in  the  air  than  most  other 
ferrous  salts  and  is  frequently  employed  in  chemical  analysis.  All  of 
the  sulfates  which  form  vitriols  yield  similar  double  salts,  not  only 
with  ammonium  sulfate  but  also  with  the  sulfates  of  potassium,  rubid- 
ium, and  caesium.  The  type  of  double  salt  represented  by  the  general 
formula  M'2SO4  •  M"SO4  •  6  H2O  therefore  includes  many  individuals, 
all  of  which  are  isomorphous. 

Ferrous  sulfide  (FeS).  Ferrous  sulfide  is  found  in  nature  as  the 
yellowish-brown  mineral  pyrrhotite,  which  nearly  always  contains  an 
excess  of  sulfur  in  solid  solution  which  may  amount  to  as  much  as 
6.5  per  cent  of  the  weight  of  the  mineral.  It  is  easily  prepared  by 
heating  iron  with  sulfur  or  by  treating  a  solution  of  a  ferrous  salt 
with  a  soluble  sulfide : 

FeCl2  +  (NH4)2S  =  FeS  +  2  NH4C1 

Prepared  in  the  latter  way  it  is  a  black  solid,  insoluble  in  water  but 
readily  soluble  even  in  very  weak  acids.  It  melts  at  about  1175°  and 
is  obtained  as  a  liquid  flux  in  some  metallurgical  processes  (p.  385). 
It  is  used  in  the  laboratory  in  the  preparation  of  hydrogen  sulfide. 

Iron  disulfide  (FeS2).  This  compound  occurs  very  abundantly  in 
nature,  especially  in  Spain.  It  is  also  found  in  the  coal  measures, 
often  forming  fossils  of  plants.  The  usual  form  is  known  as  pyrites, 
or  fool's  gold,  and  is  a  brass-yellow  mineral,  well  crystallized  in  the 
regular  system.  It  is  stable  at  moderately  high  temperatures,  whereas 
marcasite,  a  more  silvery  mineral  of  the  same  composition,  is  not 
stable  above  450°.  Little  is  known  as  to  the  structure  of  these 
compounds  or  the  valence  of  the  constituent  elements.  Pyrites  is 
mined  in  very  large  quantities  and  is  used  as  a  source  of  sulfur  in 
the  sulfuric  acid  industry  (p.  217). 


THE  IRON  FAMILY  469 

Ferrous  carbonate  (FeC03).  As  siderite,  isomorphous  with  calcite, 
ferrous  carbonate  occurs  rather  abundantly  in  nature,  often  in  large 
crystals.  Prepared  by  precipitation  it  is  a  nearly  white,  crystalline 
powder.  Like  calcite,  it  is  soluble  in  water  containing  carbon  dioxide, 
and  solutions  of  this  kind  constitute  the  chalybeate  mineral  waters. 

Ferric  compounds.  In  the  ferric  compounds  iron  acts  as  a  trivalent 
metal ;  consequently,  the  formulas  of  these  compounds  resemble  those 
of  the  corresponding  compounds  of  aluminium.  Ferric  hydroxide 
(Fe(OH)3)  is  a  very  weak  base,  and  all  the  simple  ferric  salts  are 
largely  hydrolyzed,  their  solutions  acquiring  the  reddish-brown  color 
of  the  hydroxide.  As  a  rule,  the  salts  are  not  so  well  crystallized  as 
those  of  the  ferrous  series,  and  they  present  a  greater  variety  of  color. 

Ferric  hydroxide  (Fe(OH)3).  This  insoluble  compound  is  obtained 
as  a  reddish-brown  precipitate  when  a  soluble  base  is  added  to  a  hot 
solution  of  a  ferric  salt.  If  the  solution  is  dilute  and  cold,  the  hydroxide 
remains  in  colloidal  form,  which  passes  into  the  hydrogel  when  the 
solution  is  heated.  It  forms  a  number  of  dehydration  products,  some 
of  which  are  important  ores  (p.  459).  Iron  rust  is  probably  a  mixture 
of  such  compounds.  Unlike  aluminium  hydroxide,  it  is  not  appreciably 
dissolved  by  soluble  bases. 

Ferric  oxide  (Fe203).  When  the  hydroxide  is  strongly  heated,  it 
forms  the  oxide  Fe2O3,  which  is  an  insoluble,  earthy  material  occur- 
ring in  nature  in  various  forms  of  hematite,  which  range  in  color  from 
red  to  black.  The  same  compound  is  obtained  in  burning  pyrites 
and,  when  carefully  prepared,  constitutes  the  pigment  known  as  Vene- 
tian red,  which,  owing  to  its  permanency,  is  much  used  for  paint- 
ing structures  that  are  exposed  to  the  weather,  such  as  bridges  and 
railway  cars.  This  oxide  is  also  found  in  nature  in  combination  with 
ferrous  oxide,  as  the  mineral  magnetite  (Fe3O4),  which  appears  to 
have  the  structure  of  a  spinel  (p.  444). 

Ferric  chloride  (FeCl3).  Ferric  chloride  is  obtained  in  anhydrous  form 
as  a  sublimate  by  heating  iron  in  a  current  of  chlorine.  It  is  readily 
formed  in  solution  by  the  usual  methods,  and  crystallizes  in  a  number 
of  hydrated  forms,  the  usual  one  having  the  formula  FeCl8  •  6  H2O.  It 
is  very  soluble  in  water  and  to  a  less  extent  in  other  solvents,  such 
as  alcohol  and  ether.  Its  solution  in  alcohol  constitutes  the  ordinary 
tincture  of  iron  of  the  druggist.  It  gives  characteristic  colors  with 
various  types  of  organic  compounds,  particularly  with  the  alkaloids, 
and  is  sometimes  used  as  a  reagent  in  identifying  these  compounds. 


470  GENERAL  CHEMISTRY 

Hydrolysis  of  ferric  chloride.  In  concentrated  hydrochloric  acid,  ferric  chloride 
gives  a  clear  yellow  solution,  but  in  nearly  neutral  solution  the  color  deepens  to 
a  reddish-brown  tint,  owing  to  the  hydrolysis  of  the  salt  and  the  formation  of 
the  hydroxide.  Since  the  latter  is  insoluble,  it  would  seem  reasonable  to  expect 
it  to  precipitate,  especially  as  the  dilution  increases.  It  does  not  do  so  under 
ordinary  conditions,  but  remains  as  a  colloidal  hydrosol.  It  is  only  after  long 
standing  that  the  cold  solution  deposits  the  hydrogel,  but  if  the  solution  is 
heated,  the  precipitation  is  rapid.  The  majority  of  easily  hydrolyzed  salts  exhibit 
the  same  general  conduct,  but  to  a  different  degree. 

Other  soluble  ferric  salts.  Of  the  other  soluble  ferric  salts  a  few 
deserve  special  mention.  The  sulfate  (Fe2(SO4)3)  can  be  obtained  by 
the  oxidation  of  green  vitriol,  as  a  white,  poorly  crystallized  substance. 
With  the  alkaline  sulfates  it  forms  a  series  of  violet-colored  alums, 
which,  owing  to  their  well-crystallized  condition,  are  the  ferric  salts 
most  frequently  used.  The  most  familiar  one  is  the  ferric  ammonium 
alum  (NH4Fe(SO4)2 -12H2O).  The  nitrate  is  deposited  from  concen- 
trated solutions  in  well-formed,  deliquescent  crystals  of  the  formula 
Fe(NO3)3  •  6  H2O,  which  have  the  same  violet  color  as  the  alums.  It 
is  moderately  soluble  in  dilute  nitric  acid,  but  readily  forms  super- 
saturated solutions.  The  sulfocyanate  (Fe(NCS)3)  is  a  blood-red,  solu- 
ble salt,  and  its  formation  upon  adding  a  sulfocyanate  to  a  solution 
of  unknown  composition  is  a  delicate  test  for  the  presence  of  ferric  ions. 

Insoluble  ferric  compounds.  Most  of  the  insoluble  ferric  compounds 
are  basic  salts  of  various  kinds.  The  normal  phosphate  (FePO4)  is  an 
exception  to  this,  and  is  found  in  nature  in  a  number  of  forms.  Neither 
the  sulfide  nor  the  carbonate  is  obtained  by  precipitation,  since  each  is 
completely  hydrolyzed,  as  in  the  case  of  the  corresponding  compounds 
of  aluminium. 

Oxidation  of  ferrous  salts.  When  exposed  to  the  action  of  oxidizing 
agents,  especially  in  the  presence  of  water,  ferrous  compounds  are 
readily  oxidized  to  the  corresponding  ferric  compounds.  This  is 
illustrated  in  the  case  of  the  oxide  and  hydroxide: 

4FeO  +  02  =  2Fe208 

4  Fe(OH)2  -f  2  H2O  +  O2  =  4  Fe(OH)3 

In  a  similar  way,  in  the  presence  of  sulfuric  acid  ferrous  sulfate 
is  oxidized  by  an  oxidizing  agent,  such  as  nitric  acid: 

4  FeS04  +  2  H2S04  +  O2  =  2  Fe2(SO4)3  +  2  H2O 
In  moist  air  the  oxidation  takes  place  as  follows : 

4  FeSO4  +  2  H2O  +  O2  =  4  FeSO4(OH) 


THE  IRON  FAMILY  471 

Oxidation  an  increase  in  valence.  It  will  be  noticed  that  in  these 
reactions  oxygen  is  used  up  and  the  valence  of  the  iron  is  increased 
from  2  to  3.  Any  reaction  which  increases  the  valence  of  the  metal 
of  a  salt  is  called  an  oxidation,  even  though  no  oxygen  is  involved  in 
the  process.  Thus,  ferrous  chloride  is  said  to  be  oxidized  to  ferric 
chloride  in  the  reaction  expressed  in  the  following  equation: 

2FeCl2  +  Cl2=2FeCl3 

Chlorine  is  said  to  be  an  oxidizing  agent,  since  it  effects  the  oxidation. 
Ionic  oxidation.  If  this  same  oxidation  is  represented  as  an  ionic 
reaction,  a  still  different  view  of  oxidation  is  reached.    In  this  case 
the  equation  is  as  follows  : 

2    Fe++,  C1-,  C1-   +  G1  =  2 


2 


It  will  be  seen  that  the  charge  upon  the  iron  ion  (cation)  has  increased 
from  2  to  3,  while  a  corresponding  number  of  chlorine  ions  (anions) 
have  been  formed  from  molecular  chlorine.  From  this  point  of  view 
oxidation  may  be  denned  as  a  reaction  in  which  the  charge  upon  the 
cation  has  been  increased. 

Reduction  of  ferric  compounds.  As  the  reverse  of  the  oxidation  reac- 
tions just  described,  all  ferric  compounds  may,  under  the  proper  con- 
ditions, be  reduced  to  ferrous  compounds.  This  is  illustrated  in  the 
case  of  the  oxide:  FegO§  +  H  =  2  FeO  +  H,O  : 

In  solution  many  ferric  salts  are  reduced  by  nascent  hydrogen  : 
Fe2(S04)3  +  2  H  =  2  FeSO4  +  H2SO4 

In  these  reactions  the  valence  of  the  iron  is  diminished  from  3  to  2, 
and  the  ferric  compounds  are  said  to  be  reduced  even  when  neither 
oxygen  nor  hydrogen  is  concerned  in  the  reaction.  Thus,  the  valence 
of  the  iron  may  be  diminished  by  the  action  of  a  metal  upon  a  ferric 
compound:  2FeCls  +  Fe  =  SFeCl, 

In  general,  then,  a  compound  is  said  to  be  reduced  when  the  valence 
of  the  metal  present  is  diminished. 

Ionic  reduction.  Representing  the  changes  in  the  last  two  equations 
as  ionic  reactions,  we  have  the  equations  : 


2Fe+  +  +,  3SO4--+-2H  =  2Fe++,  2H+,  3  SO4~ 
2  (Fe+  +  +,  C1-,  C1-,  C1-)  +  Fe  =  3  (Fe+  +,  Cl',  C1-) 


472  GENERAL  CHEMISTRY 

From  this  mode  of  representation  it  will  be  seen  that  the  essential 
change  in  the  reaction  is  the  decrease  in  the  charge  of  the  iron  cation 
from  3  to  2,  ami  the  reduction  of  an  electrolyte  ma^_be_defin.ed  as  a 
reaction  in  which  the  charge_on  the  cation  is  diminished.  ._. 

Oxidation  and  reduction.  If  the  earlier  definitions  of  the  terms  oxidation  and 
reduction  are  recalled  (pp.  22  and  43),  it  will  be  apparent  that  the  use  of  these 
terms  to  indicate  a  change  in  valence  involves  a  considerable  extension  of  the 
original  meaning.  In  order  that  there  may  be  no  contradiction  between  the  two 
definitions,  it  is  customary  to  assume  that  an  element  in  the  free  condition  has  a 
valence  of  zero.  In  accordance  with  this  assumption,  when  an  element  combines 
with  oxygen,  its  valence  is  increased ;  when  an  oxide  is  reduced,  the  valence  of 
the  element  is  diminished. 

Ferric  acid ;  the  ferrates.  When  fine  iron  filings  are  heated  with  potassium 
nitrate,  an  energetic  reaction  takes  place,  and  the  product  gives  a  rose-colored 
solution  with  water,  the  color  being  due  to  the  presence  of  a  salt  of  the  composi- 
tion K2FeO4.  Barium  chloride,  added  to  the  solution,  precipitates  a  difficultly 
soluble  barium  salt  of  the  formula  BaFeO4.  These  salts  are  known  as  ferrates 
and  are  analogous  in  composition  to  the  sulfates.  The  free  ferric  acid  (H2FeO4) 
has  not  been  obtained,  since  it  is  very  unstable,  decomposing  into  ferric  oxide, 
water,  and  oxygen. 

Complex  compounds  of  iron.  Iron  forms  a  large  number  of  complex 
compounds,  of  which  only  the  cyanides  will  be  described.  Iron  forms 
no  simple  cyanides,  but  a  very  large  number  of  complex  cyanides  are 
known,  of  which  potassium  ferrocyanide  and  potassium  ferricyanide 
are  the  most  important. 

Potassium  ferrocyanide  (K4Fe(CN)6).  When  nitrogenous  organic 
matter,  such  as  hoofs,  horns,  and  refuse  leather,  is  fused  with  potassium 
hydroxide  or  carbonate  and  iron  borings,  and  the  product  is  extracted 
with  water,  there  crystallizes  from  the  filtered  solution  a  beautiful 
lemon-yellow  salt,  of  the  composition  K4Fe(CN)6  •  3  H2O,  called  yellow 
prussiate  of  potash,  or  potassium  ferrocyanide.  It  is  a  true  complex, 
giving  in  solution  the  ions  4  K+  and  Fe(CN)6~ "  "  but  no  ions  of  iron. 
It  is  regarded  as  a  derivative  of  hydrogen  cyanide  (HNC),  but  in  its 
formula  the  order  of  the  symbols  N  and  C  is  usually  reversed,  this  being 
of  no  consequence,  however,  since  the  structure  is  unknown.  Unlike 
the  simple  cyanides,  it  is  not  poisonous.  The  corresponding  ferrocyanic 
acid  (H4Fe(CN)6)  can  be  obtained  as  a  white,  crystalline  precipitate  by 
treating  a  concentrated  solution  of  the  salt  with  hydrochloric  acid. 

Potassium  ferricyanide  (K3Fe(CN)6).  When  potassium  ferrocyanide 
in  solution  is  treated  with  an  oxidizing  agent,  a  greenish  solution  is 
obtained,  from  which  crystallize  garnet-red  crystals  of  the  composition 


THE  IKON  FAMILY  473 

K8Fe(CN)6,  known  as  red  prussiate  of  potash  or  potassium  ferricya- 
nide.  With  chlorine  as  the  oxidizing  agent  the  equation  is  as  follows : 

2  K4Fe(CN)6  +  C12  =  2  K8Fe(CN)6  +  2  KC1 

This  compound  in  solution  gives  the  ions  3  K+  and  Fe(CN)6  but 
no  iron  ions.  Free  ferricyanic  acid  (HgFe(CN)6)  can  be  obtained  as 
in  the  case  of  the  ferrocyanic  acid. 

Other  complex  cyanides.  Since  the  potassium  in  both  ferrocyanide  and  ferricya- 
nide  acts  as  an  ion,  it  is  replaceable  by  other  metals  through  double  decomposition. 
Most  of  the  ferrocyanides  and  ferricyanides  so  obtained  are  colloidal  bodies  in- 
soluble in  water  and  dilute  acids.  Copper  ferrocyanide  has  already  been  mentioned 
in  connection  with  its  use  as  an  osmotic  membrane.  The  compounds  obtained  by 
treating  these  two  salts  with  simple  ferrous  and  ferric  salts  are  of  especial  interest. 

With  potassium  ferrocyanide  and  ferric  chloride  the  reaction  is  represented  as 

3  K4Fe(CN)6  +  4  FeCl3  =  Fe4(Fe(CN)6)3  +  12  KC1 

The  complex  product  of  this  reaction  is  an  indigo-blue  precipitate,  known  as 
ferric  ferrocyanide,  or  Prussian  blue,  and  is  used  largely  as  a  pigment. 
With  potassium  f erricyanide  and  ferrous  sulf ate  the  equation  is 

2  K3Fe(CN)6  +  3  FeSO4  =  Fe3(Fe(CN)6)2  +  3  K2SO4 

The  resulting  complex  is  of  much  the  same  color  as  Prussian  blue  and  is  known 
as  ferrous  ferricyanide,  or  TurnbuJPs  blue.  The  formation  of  these  two  precipi- 
tates affords  a  method  of  distinguishing  between  ferrous  and  ferric  salts. 

Blue  printing.  When  a  ferric  salt  and  potassium  ferricyanide  are  brought  to- 
gether in  solution,  no  precipitate  forms,  though  the  solution  acquires  a  yellowish 
color.  On  exposure  to  the  sunlight  the  ferric  salt  undergoes  a  partial  reduction  to 
ferrous  salt,  and  a  blue  precipitate  forms.  Advantage  is  taken  of  these  facts  in 
the  process  of  blue  printing.  A  sensitive  paper  is  prepared  by  soaking  paper  in 
a  solution  of  potassium  ferricyanide  and  a  ferric  salt  (ferric  ammonium  citrate 
is  generally  used)  and  drying  it  in  a  dark  place.  When  a  black  drawing  on  trac- 
ing cloth  is  placed  upon  such  a  sensitive  paper  and  the  two  are  exposed  to  the 
sunlight,  the  sensitive  paper  (except  where  it  is  protected  by  the  black  lines)  turns 
a  brownish  color.  It  is  then  thoroughly  washed  with  water,  to  remove  the  soluble 
salts,  the  portions  acted  upon  by  the  light  turning  blue,  while  the  unaffected 
portions  are  left  white.  Both  Prussian  blue  and  Turnbull's  blue  are  decomposed 
and  thus  decolorized  by  soluble  bases,  so  that  a  solution  of  sodium  hydroxide 
can  be  used  as  an  ink  for  white  lettering  on  a  blue  print. 

COBALT 

Occurrence.  Most  minerals  containing  cobalt  are  strongly  suggestive 
of  the  presence  of  a  heavy  metal,  yet  the  older  metallurgists  were 
unable  to  smelt  them  and  obtain  this  metal.  For  this  reason  they 
named  the  metal  kobold,  meaning  "goblin,"  and  this  gave  us  our 
name  cobalt.  The  metal  was  finally  obtained  by  the  Swedish  chemist 


474  GENERAL  CHEMISTKY 

Brandt  in  1735.  Cobalt  usually  occurs  in  combination  with  arsenic 
and  sulfur,  in  complex  minerals  which  also  contain  silver,  iron,  nickel, 
and  copper.  The  simplest  of  these  is  cobaltite  (CoAsS).  Such  min- 
erals are  found  sparingly  in  many  localities,  but  the  richest  deposits 
are  those  of  New  Caledonia  and  Ontario.  The  latter  are  very  rich  in 
silver,  and  cobalt  is  worked  up  as  a  by-product,  together  with  nickel 
and  arsenic  trioxide. 

Metallurgy  and  properties.  The  metallurgy  of  the  metal  is  very  com- 
plicated, since  it  is  difficult  to  separate  cobalt  and  nickel.  The  pure 
metal  is  best  prepared  by  the  Goldschmidt  process.  It  is  a  malle- 
able, magnetic,  silvery  metal,  which  soon  takes  on  a  reddish  tint  upon 
exposure  to  the  air.  It  melts  at  1490°  and  has  a  density  of  8.6.  The 
metal  itself  has  no  uses  in  the  industries,  but  its  alloys  are  attracting 
attention,  particularly  one  which  it  forms  with  chromium  (p.  520).  It 
can  be  plated  upon  other  metals  like  nickel,  which  it  greatly  resembles. 

Compounds.  Cobalt  comes  into  the  market  very  largely  in  the  form 
of  the  black,  cobaltous  oxide  CoO,  and  in  addition  to  this  it  forms 
two  other  oxides,  Co2O3  and  Co3O4,  corresponding  to  those  of  iron. 
The  cobaltous  oxide  gives  rise  to  a  series  of  salts  similar  in  formulas 
to  the  ferrous  salts.  In  anhydrous  form  they  are  blue,  but  when  hy- 
drated  they  are  usually  cherry  red.  The  simple  salts  derived  from 
cobaltic  oxide,  corresponding  to  the  ferric  salts,  are  very  unstable, 
and  few  are  well  known.  There  are,  however,  a  great  many  complex 
cobaltic  salts. 

Cobaltous  oxide  (CoO).  This  is  a  black  powder  used  in  making  other 
cobalt  compounds  and  in  making  blue  glass  and  blue  decorations  on 
china.  When  used  as  an  ingredient  in  glasses,  glazes,  and  enamels,  it 
forms  a  blue  silicate  which  has  intense  coloring  properties.  Sometimes 
the  ground  blue  glass  called  smalt  is  used  instead  of  the  oxide,  as 
well  as  for  a  pigment. 

Salts  of  cobalt.  The  hydrated  nitrate  (Co(NO3)2  •  6  H2O)  and  the 
chloride  (CoCl2  •  6  H2O)  are  -the  salts  most  frequently  employed  in 
the  laboratory;  both  these  salts  are  cherry  red  in  color.  The  sulfide 
(CoS)  is  formed  as  a  black  precipitate  when  a  cobalt  salt  is  treated 
with  ammonium  sulfide.  It  dissolves  slowly  in  dilute  acids  but  much 
more  rapidly  in  concentrated  ones. 

When  sodium  nitrite,  potassium  chloride,  and  a  salt  of  cobalt  are  brought  into 
solution  and  treated  with  dilute  acetic  acid,  a  complicated  reaction  takes  place, 
in  which  the  cobalt  is  oxidized  to  the  trivalent  state  by  the  nitrous  acid  liberated 


THE  IRON  FAMILY  475 

in  the  reaction,  and  an  insoluble  yellow  salt  is  precipitated,  the  formula  of  which 
is  K3Co(NO2)6.  It  is  known  as  potassium  cobaltinitrite,  or  Fischer's  salt,  and  its 
formation  is  employed  as  a  test  both  for  cobalt  and  for  potassium. 

Cobaltammines.  When  a  salt  of  cobalt  is  treated  with  ammonia,  the 
pale  blue  hydroxide  Co(OH)2  is  at  first  precipitated.  This  dissolves 
in  excess  of  ammonia,  absorbs  oxygen  from  the  air,  and  is  oxidized 
to  the  trivalent  state.  From  solutions  prepared  in  this  way  highly 
colored  complex  salts  are  obtained,  the  composition  of  which  depends 
upon  the  salt  of  cobalt  used,  other  compounds  which  may  be  present, 
and  the  conditions  of  the  experiment.  As  a  class  these  salts  are  called 
the  cobaltammines,  and  a  great  number  of  them  have  been  prepared. 
The  best-known  type  is  represented  by  the  formula  Co(NH3)6Cl3,  which 
is  named  hexammino-cobalti-chloride.  These  compounds  have  many 
interesting  peculiarities,  but  a  study  of  them  would  take  us  too  far. 

NICKEL 

Occurrence.  The  early  metallurgists  were  acquainted  with  certain 
minerals  of  high  metallic  luster  which  strongly  resembled  the  ores  of 
copper,  but  from  which  they  could  not  extract  that  metal,  and  which 
they  accordingly  named  kupfemickel,  or  false  copper.  The  labors  of 
Cronstedt  and  Bergman  in  Sweden,  toward  the  close  of  the  eighteenth 
century,  resulted  in  the  isolation  of  the  new  metal  and  its  clear 
differentiation  from  cobalt. 

Nickel  is  almost  always  associated  with  cobalt  in  nature.  Like  the 
latter  element,  it  occurs  in  combination  with  sulfur  and  arsenic  and  as- 
sociated with  copper,  silver,  and  iron.  Most  of  its  ores  are  very  com- 
plex, and  it  was  formerly  obtained,  chiefly  as  a  by-product,  in  the 
metallurgy  of  copper  and  silver.  The  rich  ores  of  New  Caledonia 
and  Ontario  are  now  worked  primarily  for  nickel,  the  chief  mineral 
being  garnierite  (2(Ni,  Mg)gSi4O13  •  3  H2O),  the  formula  indicating 
that  a  variable  quantity  of  nickel  may  replace  the  magnesium  of  the 
mineral. 

Metallurgy.  The  extraction  of  nickel  from  its  ores  is  a  very  com- 
plicated process.  The  essential  features  are  the  roasting  of  the  ores 
until  the  nickel  and  cobalt  are  left  as  oxides,  and  the  subsequent  re- 
duction of  the  oxides  with  carbon.  The  separation  of  the  cobalt  and 
nickel  presents  many  difficulties.  Several  million  pounds  of  nickel 
sulfate  are  recovered  annually  in  the  United  States  from  the  copper 
sulfate  baths  used  in  the  refining  of  copper. 


476  GENERAL  CHEMISTRY 

Properties  and  uses.  Nickel  is  a  silvery  metal  capable  of  a  very 
high  polish.  It  is  very  hard,  but  is  quite  malleable.  It  can  be  welded 
on  iron  and  the  two  rolled  into  sheets  for  making  various  kitchen 
utensils.  Like  iron  and  cobalt,  it  is  magnetic.  Its  density  is  8.9  and 
it  melts  at  1452°.  It  is  not  attacked  by  melted  alkalies,  and  nickel 
crucibles  are  often  employed  in  the  laboratory  for  alkali  fusions.  The 
nonoxidizing  acids  evolve  hydrogen  with  nickel  very  slowly,  but  nitric 
acid  dissolves  it  readily.  It  is  used  chiefly  as  a  constituent  of  alloys, 
such  as  nickel-steel  and  coinage  metals  (see  table,  p.  480),  and  for 
plating  upon  other  metals. 

Oxides  of  nickel.  Nickel  forms  three  well-known  oxides  of  the 
formulas  NiO,  Ni2O3,  and  NigO4,  corresponding  to  those  of  iron  and 
cobalt.  It  also  forms  several  others  which  are  less  well  denned.  Of 
these,  nickelous  oxide  (NiO)  is  the  only  one  which  gives  rise  to  a 
series  of  simple  salts,  corresponding  to  the  ferrous  salts.  When  an- 
hydrous, these  are  usually  yellow,  and  when  hydrated,  some  shade  of 
deep  green.  Only  a  few  of  these  require  description. 

Salts  of  nickel.  Nickel  sulfide  (NiS),  as  prepared  by  precipitation,  is  a 
black,  amorphous  powder,  insoluble  in  water  and  dilute  acids,  but  easily 
soluble  in  more  concentrated  acids.  Nickel  chloride  (NiCl2  •  6  H2O), 
nickel  nitrate  (Ni(NO3)2  •  6  H2O),  and  nickel  sulfate  (NiSO~4  •  7  HgO) 
are  the  most  familiar  simple  salts.  The  sulfate  also  forms  crystals  of 
the  composition  (NH4)2SO4  •  NiSO4  •  6  H2O,  corresponding  to  Mohr's 
salt.  It  is  the  salt  of  nickel  employed  as  the  electrolyte  in  nickel  plat- 
ing, a  piece  of  pure  nickel  being  used  as  the  anode  and  the  object  to 
be  plated  as  the  cathode.  There  are  many  complex  salts  of  nickel. 

Nickel  carbonyl  (Ni(CO)J.  When  carbon  monoxide  is  passed  over 
metallic  nickel  at  a  temperature  between  30°  and  50°,  the  two  unite  to 
form  a  compound  of  the  formula  Ni(CO)4,  known  as  nickel  carbonyl. 
It  is  a  colorless  liquid  boiling  at  43.2°  and  freezing  at  —  25°  to  color- 
less, needle-shaped  crystals.  When  the  vapor  of  the  compound  is 
passed  through  a  tube  heated  to  above  100°,  the  compound  dissociates 
into  the  metal  and  carbon  monoxide.  Advantage  is  taken  of  this 
reaction  in  the  Mond  process  for  purifying  nickel. 

Carbonyls  of  other  metals.  Cobalt  forms  two  carbonyls,  Co(CO)3  and  (CoCO)4, 
but  both  of  these  are  solids  and  are  formed  only  under  pressure.  Iron  forms  three  ; 
namely,  Fe(CO)4,  Fe(CO)5,  and  Fe2(CO)9.  Of  these  the  pentacarbonyl  is  a  yellow 
liquid  boiling  at  103°.  It  sometimes  forms  in  gas  pipes  and,  upon  its  combustion, 
clogs  the  burner  with  iron  oxide  or  injures  the  gas  mantle  by  a  deposit  upon  it. 


CHAPTER  XXXV 

COPPER;  MERCURY;  SILVER 


SYMBOL 

ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

BOILING 
POINT 

CONDUC- 
TIVITY 

Copper  .     . 

Cu 

63.57 

8.93 

1082.6° 

2310.° 

0.561 

Mercury     . 

Hg 

200.6 

13.56 

-  38.8° 

1    356.7° 

0.104 

Silver    .     . 

Ag 

107.88 

10.5 

960>° 

1955.° 

0.614 

General.  Although  these  three  elements  do  not  form  a  periodic 
family,  copper  and  silver  belonging  to  one  family  and  mercury  to 
another,  they  have  much  in  common,  and  it  is  convenient  to  describe 
them  together. 

Occurrence.  In  nature  they  are  all  found  to  some  extent  as  native 
metals,  which  indicates  that  they  are  rather  inactive  chemically.  For 
the  most  part  they  occur  in  combination  with  sulfur,  as  either  simple 
or  complex  sulfides.  The  reactions  involved  in  their  metallurgy  are 
therefore  not  unlike,  and  lead  to  methods  different  from  those  so 
far  described. 

Properties.  All  three  are  metals  of  rather  high  density  and  are  good 
conductors  of  electricity,  silver  and  copper  surpassing  all  other  metals 
in  this  respect.  These  two  are  also  exceptionally  ductile  and  malleable. 
In  liquid  condition  they  mix  well  with  each  other  and  also  with  other 
metals,  and  so  form  numerous  alloys. 

Chemical  conduct.  In  chemical  conduct  they  are  rather  inactive, 
especially  toward  oxygen,  their  oxides  being  very  easily  reduced. 
Toward  sulfur  they  are  much  more  active,  and  their  sulfides  are 
more  stable  than  their  oxides.  They  do  not  displace  hydrogen  from 
acids  and  water,  since  they  are  below  it  in  the  electromotive  series. 
Oxidizing  acids,  such  as  nitric  and  sulfuric,  convert  them  into  the 
corresponding  salts.  Silver  forms  but  one  series  of  salts,  and  in  this 
the  metal  is  univalent.  Copper  and  mercury  each  form  two  series. 
In  the  one,  known  as  the  cuprous  and  mercurous  series,  they  are 
univalent,  while  in  the  other,  known  as  the  cupric  and  mercuric 
series,  they  are  divalent. 

477 


4T8 


GENERAL  CHEMISTRY 


COPPER 


Occurrence.  Metallic  copper  has  been  known  from  the  earliest  times 
and  was  probably  the  first  metal  to  come  into  any  considerable  use. 
This  is  explained  by  its  native  occurrence  and  by  the  ease  with  which 
its  oxygen  compounds  are  reduced.  It  owes  its  name  (from  the  Latin 
word  cuprum)  to  the  fact  that  the  Romans  obtained  it  from  the  island 
of  Cyprus. 

Large  quantities  of  native  copper,  in  a  state  approximating  purity, 
are  found  in  the  northern  peninsula  of  Michigan.  Smaller  deposits 
are  also  found  in  a  number  of  other  localities.  In  combination  it  is 
found  in  a  wide  variety  of  forms,  especially  as  sulfides,  oxides,  car- 
bonates, and  silicates.  In  small  quantities  it  is  present  in  the  great 
majority  of  sulfide  minerals,  and  in  traces  it  is  very  widely  distributed 
in  nature.  To  some  extent  it  is  absorbed  by  plants  growing  in  copper- 
bearing  regions,  and  it  is  assimilated  by  some  animals,  as  is  shown  by 
the  fact  that  it  occurs  in  the  blood  of  the  cuttlefish  and  in  the  feathers 
of  some  birds. 

Ores  of  copper.  The  table  which  follows  gives  a  list  of  the  most  impor- 
tant ores,  chalcopyrite  and  native  copper  being  by  far  the  most  valuable. 


SULFIDE  OKES 

OXYGEN 

ORES 

C  halcopyrite 

CuFeS 

Cuprits 

On  O 

Chalcocits 

Cu  S 

M^laconite 

CuO 

Bornite      

.  CuQFeS 

Malachite 

CuCO  •  Cu(OH) 

The  most  important  copper-producing  states  in  the  United  States 
are  Arizona,  Montana,  Michigan,  and  Utah.  Nearly  all  civilized  coun- 
tries produce  some  copper,  but  the  United  States  produces  more  than 
one  half  of  the  world's  supply. 

Metallurgy.  Ores  containing  little  or  no  sulfur  are  easily  reduced, 
it  being  only  necessary  to  heat  them  in  a  suitable  furnace  together 
with  coke  and  an  appropriate  flux.  The  slag  resulting  from  this 
process,  however,  carries  away  considerable  copper,  and  on  this  ac- 
count it  is  more  economical  to  mix  the  ores  with  others  containing 
sulfur  and  to  use  a  different  method.  The  sulfide  ores  always  contain 
iron  and  usually  a  little  silver  and  gold.  The  problem,  therefore,  is  to 
separate  the  copper,  together  with  the  silver  and  gold,  from  the  sulfur 
and  iron.  In  a  general  way  this  is  accomplished  in  a  series  of  opera- 
tions by  heating  the  ore,  together  with  silica,  in  a  regulated  current 


COPPER;  MEKCURY;  SILVER  479 

of  air.  The  sulfur  burns  to  sulfur  dioxide,  which  then  escapes.  The 
iron  is  first  converted  into  oxide,  which  then  combines  with  silica  to 
form  a  slag  of  iron  silicate,  while  the  copper,  silver,  and  gold  are  left 
as  a  metallic  alloy. 

The  details  of  the  process  are  complicated  and  vary  greatly  in  different 
plants.  In  the  United  States  the  process  in  widest  use  resembles  in  many  respects 
the  one  by  which  steel  is  made,  being  based  upon  the  use  of  a  blast  furnace  and 
a  Bessemer  converter.  When  the  ore  is  too  fine  to  be  treated  in  a  blast  furnace, 
a  mixture  of  oxide  and  sulfide  ores  is  first  roasted  in  a  reverberatory  furnace, 
care  being  taken  to  leave  enough  sulfur  to  combine  with  all  of  the  copper  and  a 
part  of  the  iron.  The  ore  so  treated,  or  coarse  sulfide  ores  which  require  no  pre- 
liminary roasting,  together  with  a  flux  rich  in  silica,  is  charged  into  a  blast  fur- 
nace known  as  the  matte  furnace.  In  this  the  iron  oxide  combines  with  silica 
to  form  slag,  while  the  copper,  copper  sulfide,  iron  sulfide,  silver,  and  gold  melt 
to  a  heavy  liquid  called  matte.  By  repeating  the  process,  if  necessary,  the  matte 
is  brought  to  a  content  of  as  much  as  50  per  cent  copper.  The  hot  matte  is  then 
poured  into  a  converter,  closely  resembling  a  Bessemer  converter,  which  holds 
from  6  to  10  tons.  A  suitable  quantity  of  silica  is  also  added,  and  air  is  blown 
through  the  liquid.  The  sulfur  acts  as  fuel,  burning  to  form  the  dioxide,  while 
the  iron  oxide  produced  passes  into  the  slag  as  silicate.  This  slag  floats  upon  the 
surface  of  the  melted  copper  and  is  run  off  and  returned  to  the  matte  furnace, 
while  the  nearly  pure  copper  is  poured  out  into  molds.  The  product  is  called 
blister  copper  and  may  have  a  purity  of  as  much  as  98  per  cent. 

Refining  of  copper.  Many  of  the  uses  for  which  copper  is  employed 
require  a  very  pure  metal,  and  for  these  purposes  blister  copper  must 
be  refined.  This  is  accomplished  by  electrolysis. 

The  copper  from  the  converter  is  cast  into  anode  plates  weighing  upward  of 
300  pounds.  These  are  suspended  in  tanks  containing  a  solution  of  copper  sul- 
fate  as  electrolyte,  each  anode  plate  being  arranged  opposite  to  a  cathode  made 
of  a  thin  sheet  of  pure  copper.  The  current,  in  passing  through  the  cell,  dissolves 
copper  from  the  anode  and  deposits  it  upon  the  cathode  in  very  pure  form,  the 
insoluble  impurities  collecting  on  the  bottom  of  the  tank  as  a  mud.  The  cathode 
copper,  while  pure,  is  porous  and  is  melted  and  cast  into  compact  ingots.  The 
electrolytic  mud  contains  the  gold  and  silver  which  was  in  the  blister  copper  and 
is  worked  over  to  obtain  these  precious  metals.  It  often  contains  tellurium  as  well, 
which  at  present  has  no  commercial  value. 

Properties  of  copper.  Copper  is  a  heavy  metal  of  characteristic  ruddy 
color,  whose  density  averages  about  8.934.  It  melts  at  1082.6°  and  boils 
at  2310°.  It  is  rather  soft  and  is  very  ductile,  malleable,  and  flexible, 
yet  tough  and  fairly  strong.  As  an  electrical  conductor  it  is  second 
only  to  silver.  Its  properties,  notably  its  electrical  conductivity,  are 
much  altered  by  impurities,  especially  by  the  presence  of  oxides  and 
sulfides.  It  is  not  attacked  by  nonoxidizing  acids,  unless  oxygen  is 


480  GENERAL  CHEMISTRY 

present,  nor  by  fused  alkalies,  but  oxidizing  acids  convert  it  into  the 
corresponding  salts.  In  the  presence  of  air  most  acids  slowly  act  upon 
it ;  even  carbon  dioxide  in  moist  air  gradually  covers  its  surface  with 
a  greenish  coating  of  a  basic  carbonate.  When  heated  in  the  air,  it  is 
oxidized  to  black  copper  oxide  (CuO).  Sulfur  and  the  halogens  attack 
it  with  much  more  energy  than  does  oxygen. 

Uses  of  copper.  Copper  is  extensively  used  in  electrical  construction, 
as  a  constituent  of  alloys,  for  roofing,  for  sheathing  the  bottoms  of 
ships,  for  coinage,  and  for  many  minor  purposes.  The  following  table 
gives  the  names  and  approximate  composition  of  some  of  its  most 
important  alloys. 

Aluminium  bronze     ....  90-98%  copper,  2-10%  aluminium  ) 

Brass 63-73%  copper,  27-37%  zinc 

Bronze 70-95%  copper,  1-25%  zinc,  1-18%  tin 

German  silver 56-60%  copper,  20%  zinc,  20-25%  nickel 

Gun  metal 90%  copper,  10%  tin 

Gold  coin 10%  copper,  90%  gold 

Silver  coin 10%  copper,  90%  silver 

Nickel  coin 75%  copper,  25%  nickel 

Electrotyping.  Books  are  often  printed  from  electrotype  plates,  which  are  pre- 
pared as  follows :  The  face  of  the  type  is  covered  with  wax,  and  this  is  firmly 
pressed  down  until  a  clear  impression  is  obtained.  The  impressed  side  of  the 
wax  is  coated  with  graphite,  and  this  is  made  the  cathode  in  an  electrolytic  cell 
containing  a  copper  salt  in  solution.  The  copper  is  deposited  as  a  thin  sheet 
upon  the  letters  in  wax  and,  when  detached,  is  a  perfect  copy  of  the  type,  the 
under  part  of  the  letters  being  hollow.  The  sheet  is  strengthened  by  pouring 
on  the  under  surface  a  suitable  amount  of  commercial  lead.  The  sheet  so 
strengthened  is  then  used  in  printing. 

Simple  compounds  of  copper.  Copper  forms  two  series  of  simple  salts 
derived  from  the  oxides  Cu2O  (cuprous)  and  CuO  (cupric).  Under 
ordinary  conditions,  in  the  presence  of  moisture  and  air,  the  cupric 
salts  are  much  the  more  stable,  while  at  high  temperatures  the  cuprous 
salts  are  the  stable  form.  In  aqueous  solutions  cupric  compounds  are 
blue,  while  cuprous  compounds  are  colorless.  Both  cuprous  hydroxide 
(CuOH)  and  cupric  hydroxide  (Cu(OH)2)  are  rather  weak  bases,  and 
their  salts  are  somewhat  hydrolyzed  in  solution,  giving  the  solution 
an  acid  reaction.  They  also  form  many  basic  salts,  which  are  very 
sparingly  soluble.  All  copper  salts  are  more  or  less  poisonous,  es- 
pecially to  lower  forms  of  life,  and  a  number  of  them  are  used  as 
insecticides.  Copper  and  its  salts  have  a  catalytic  action  in  a  great 
many  chemical  reactions,  an  example  being  Deacon's  process  (p.  249). 


COPPER;  MERCURY;  SILVER,  481 

Cuprous  compounds.  Simple  cuprous  salts  are  rarely  derived  from 
oxygen  acids,  but  are  represented  by  such  compounds  as  the  oxide, 
sulfide,  the  halides,  and  the  cyanide.  They  are  prepared  in  two  gen- 
eral ways  : 

1.  By  heating  a  solution  of  a  cupric  salt  with  a  reducing  agent.  The 
simplest  reactions  of  this  kind  are  those  in  which  metallic  copper  is 
employed  as  the  reducing  agent.  The  preparation  of  cuprous  chloride 
is  an  example  : 


2.  By  the  dissociation  of  cupric  compounds.  At  higher  temperatures 
most  binary  cupric  compounds  dissociate  to  form  cuprous  compounds  : 

4  CuO  =  2  Cu2O  +  O2 
2  CuS  =  Cu2S  +  S 

Cuprous  compounds  are  nearly  all  very  sparingly  soluble  in  water  and 
are  strong  reducing  agents  tending  to  pass  into  cupric  salts. 

Cuprous  oxide  (Cu20).  Cuprous  oxide  is  found  in  nature  as  a  deep 
red  mineral  called  ruby  copper,  or  cuprite.  It  is  formed  as  a  precipi- 
tate when  cupric  compounds  are  heated  with  a  reducing  agent  in  alka- 
line solution.  Cupric  hydroxide  is  first  formed,  and  this  is  then  reduced 
as  follows  : 

2  Cu(OH)2  =  H2O  +  O  +  2  CuOH  —  >-  Cu2O  +  H2O 

Certain  sugars  in  solution  effect  this  reduction,  and  the  formation  of 
cuprous  oxide  is  often  employed  as  a  test  for  these. 

Cuprous  sulfide  (Cu2S).  This  compound  is  found  in  nature  as  a  brass- 
yellow  mineral  called  chalcocite.  It  is  formed  when  cupric  sulfide  is 
heated  in  the  absence  of  air,  preferably  in  a  current  of  hydrogen  : 


Cuprous  chloride  (CuCl).  Cuprous  chloride  is  a  snow-white,  crystal- 
line solid,  almost  insoluble  in  water.  It  is  most  easily  prepared  by 
heating  a  solution  of  cupric  chloride  with  copper  turnings,  or  by  pass- 
ing a  current  of  sulfur  dioxide  into  a  hot  solution  of  a  cupric  salt  and 
hydrochloric  acid  : 

2  CuSO4  +  2  HC1  +  2  H2O  +  SO2  =  2  CuCl  +  3  H2SO4 

While  practically  insoluble  in  water,  it  is  readily  soluble  in  concen- 
trated hydrochloric  acid  and  in  ammonia  water,  in  both  cases  forming 
complex  compounds.  These  solutions  absorb  many  gases,  especially 


482  GENERAL   CHEMISTRY 

oxygen,  carbon  monoxide,  and  acetylene.  With  acetylene  a  chocolate- 
colored  precipitate  of  cuprous  acetylide  is  thrown  down : 

C2H2+  2  CuCl  =  Cu2C2  +  2  HC1 

With  carbon  monoxide  an  addition  product  is  formed : 
CuCl  +  CO  +  H2O  -  CuCl  •  CO  •  H2O 

The  common  method  used  for  determining  the  percentages  of  these 
two  gases  present  in  gas  mixtures  is  usually  based  upon  these 
reactions. 

Cuprous  bromide  (CuBr)  and  cuprous  iodide  (Cul).  These  com- 
pounds have  properties  quite  similar  to  those  of  the  chloride,  but  they 
are  more  readily  formed  from  the  corresponding  cupric  salts.  In  the 
case  of  the  iodide  it  is  not  necessary  to  employ  a  reducing  agent,  since 
cupric  iodide  spontaneously  decomposes  into  the  cuprous  salt  and  free 
iodine.  Being  insoluble,  it  precipitates  when  a  solution  of  an  iodide 
is  added  to  any  simple  cupric  salt : 

2  CuSO4  +  4  KI  =  2  K2SO4  +  2  Cul  +  I2 

Since  iodine  can  be  .very  accurately  estimated  by  means  of  sodium 
thiosulfate  (p.  400),  the  quantity  of  copper  present  in  a  solution  can 
be  determined  with  precision  by  measuring  the  iodine  set  free,  and 
this  reaction  is  much  used  in  the  commercial  estimation  of  copper. 

Cuprous  cyanide  (CuNC).  This  white,  insoluble  salt  is  formed  by  a 
reaction  quite  analogous  to  the  one  in  which  cuprous  iodide  is  obtained, 
since  cupric  cyanide  spontaneously  decomposes  into  the  cuprous  salt 
and  cyanogen: 

2  CuSO4  +  4  KNC  =  2  K2SO4  +  2  CuNC  +  C2N2 

Cupric  salts.  The  cupric  salts  are  the  familiar  salts  of  copper.  In 
most  cases  they  are  obtained  by  the  usual  methods,  and  a  great  variety 
of  them  are  known,  both  normal  and  basic.  In  the  solid  state  they  are 
usually  'blue,  green,  or  yellow ;  in  dilute  solution  they  are  all  blue. 
Some  of  them  can  be  employed  as  mild  oxidizing  agents,  being 
converted  into  cuprous  compounds.  Only  a  few  require  detailed 
description. 

Cupric  oxide  (CuO).  The  black  oxide  of  copper  is  usually  obtained 
commercially  by  heating  copper  powder  or  turnings  in  the  air.  It  is 
insoluble  in  water,  but  is  readily  soluble  in  acids,  yielding  the  corre- 
sponding salts.  Owing  to  the  ease  with  which  it  gives  up  its  oxygen,  it 


COPPER;  MERCURY;  SILVER      .  483 

is  a  good  oxidizing  agent  and  finds  extensive  use  in  the  laboratory  for 
such  operations  as  the  quantitative  oxidation  of  carbon  compounds. 
Industrially  it  is  used  on  a  large  scale  in  the  refining  of  petroleum 
for  the  removal  of  sulfur  from  the  oil.  It  is  regenerated  from  the 
resulting  sulfur  compounds  by  roasting  in  air,  and  is  used  over  again. 
Cupric  hydroxide  (Cu(OH)2).  The  insoluble  hydroxide  results  as 
a  pale  blue,  colloidal  precipitate  when  .any  cupric  salt  is  brought 
together  with  a  solution  of  a  metallic  base: 

CuS04  +  2  KOH  =  K2S04  +  Cu(OH)2 

When  the  mixture  is  heated,  the  hydroxide  is  converted  into  oxide  : 
Cu(OH)2  =  CuO  +  H20 

Cupric  sulfide  (CuS).  Cupric  sulfide  results  as  a  brownish-black  pre- 
cipitate when  a  solution  of  a  cupric  salt  is  treated  with  hydrogen 

CuSO4  +  H2S  =  CuS  +  H2SO4 

It  is  insoluble  both  in  water  and  in  dilute  acids.  It  is  also  nearly  in- 
soluble in  a  solution  of  ammonium  sulfide,  which  distinguishes  it  from 
some  other  sulfides. 

Cupric  sulfate  (CuSOJ.  In  its  industrial  uses  cupric  sulfate  is  the 
most  important  of  the  salts  of  copper.  Under  ordinary  conditions  it 
crystallizes  from  solution  in  blue  triclinic  crystals,  often  of  very  large 
size,  having  the  composition  CuSO4  •  5  H2O  and  known  as  blue  vitriol, 
or  bluestone.  A  number  of  other  hydrates  are  known,  all  of  which, 
when  strongly  heated,  yield  the  anhydrous  salt,  which  is  white.  In  the 
industries  it  is  obtained  in  the  refining  of  silver  (p.  495)  and  by  the 
oxidation  of  pyrite  containing  copper  : 


Prepared  in  these  ways  it  contains  a  considerable  percentage  of  ferrous 
sulfate  as  an  impurity.  It  is  used  as  a  source  of  copper  in  the  manu- 
facture of  other  copper  salts,  as  an  electrolyte  in  copper  refining,  in 
electrotyping,  and  in  batteries,  and  for  the  treatment  of  hoof  diseases, 
particularly  in  sheep.  A  solution  containing  cupric  sulfate,  potassium 
sodium  tartrate,  and  sodium  hydroxide  is  known  as  Fehling's  solution 
and  is  used  in  the  determination  of  certain  sugars  (p.  481).  The  ordi- 
nary insecticide  known  as  Bordeaux  mixture  is  made  by  adding  cal- 
cium hydroxide  to  a  cold  solution  of  copper  sulfate.  Many  lower 
organisms,  particularly  those  known  as  algae,  are  destroyed  by  even 


484  .  GENERAL  CHEMISTRY 

very  small  traces  of  soluble  copper  salts,  and  copper  sulfate  is  some- 
times added  to  the  water  supply  of  cities  to  kill  the  algae,  whose  growth 
imparts  an  unpleasant  taste  and  odor  to  the  water. 

Cupric  carbonate.  The  normal  carbonate  of  copper  is  not  Jmojgn,  but 
there  are  a  number  of  basic  carbonates,  the  chief  of  which  have  the 
formulas  CuCO3  •  Cu(OH)2  (which  occurs  in  nature  as  the  green  mal- 
achite) and  2  CuCO3  •  Cu(OH)2  (which  is  the  blue  mineral  azurite). 

Other  cupric  salts.  Among  the  other  cupric  salts  frequently  used  in 
the  laboratory  are  the  following,  most  of  which  form  other  hydrates 
in  addition  to  those  given : 

Cupric  nitrate  (Cu(NO3)2  •  6  H2O) :  blue,  deliquescent  crystals. 

Cupric  chloride  (CuCl2*  2  H2O) :  light  blue,  pearly  scales,  or  needles. 

Cupric  bromide  (CuBr2):  brownish-purple  crystals  resembling  iodine. 

Cupric  acetate  (Cu(C2HgO2)2-  H2O)  :  a  blue,  easily  crystallized  salt. 

Complex  salts  of  copper.  Like  nearly  all  the  metals  of  high  density, 
copper  forms  a  great  many  double  and  complex  salts,  a  study  of  which 
would  take  us  too  far.  Only  two  general  classes  will  be  mentioned 
here;  these  are  of  importance  in  themselves  and  represent  classes 
which  will  frequently  recur  with  other  metals. 

1.  Ammonia  compounds.  When  cupric  sulfate  is  treated  with  aqua 
ammonia,  the  insoluble  hydroxide  is  at  first  precipitated,  as  would  be 
expected.  Continued  addition  of  ammonia  causes  the  precipitate  to 
dissolve,  forming  an  intensely  blue-purple  solution.  From  this  solu- 
tion, under  favorable  conditions,  there  crystallizes  a  solid  of  the  same 
intense  color,  which  has  the  formula  Cu(NH3)4SO4  •  H2O.  The  great 
majority  of  cupric  salts  yield  similar  compounds,  with  excess  of  am- 
monia, and  all  have  the  same  intense  color,  quite  unlike  the  pale  blue 
of  simple  copper  salts.  This  appears  to  be  due  to  the  fact  that  the 
copper  combines  with  ammonia  to  form  the  complex  ion  Cu(NHg)4++, 
the  sulfate  ionizing  according  to  the  following  equation : 

Cu(NH8)4S04  *=±  Cu(NH3)4++  +  SO,-  - 

The  dissolving  of  a  precipitate  by  continued  addition  of  the  reagent 
which  produced  it  may  'Usually  be  taken  to  indicate  the  forma- 
tion of  a  complex  ion  of  this  general  kind.  Cupriammonium  chloride 
(Cu(NH3)4Cl2)  is  formed  in  the  same  way  from  copper  chloride. 

Under  similar  circumstances  cuprous  compounds  yield  colorless 
compounds  in  which  there  are  usually  two  ammonia  groups.  Thus, 
cuproammonium  chloride  has  the  formula  Cu(NH3)2Cl. 


COPPER;  MERCURY;  SILVER 


485 


2.  Complex  cyanides.  We  have  seen  that  when  a  cupric  salt  and  a 
soluble  cyanide  are  brought  together  in  solution,  insoluble  cuprous 
cyanide  (CuNC)  is  formed.  If  an  excess  of  the  cyanide  is  added,  the 
precipitate  dissolves,  and  the  resulting  solution  is  colorless.  When 
treated  with  the  reagents  which  usually  precipitate  copper  compounds, 
—  for  example,  with  hydrogen  sulfide,  —  it  gives  no  precipitate,  nor 
does  it  turn  blue-purple  with  ammonia.  These  facts  point  to  the 
formation  of  a  complex  ion,  and  experiment  shows  that  the  product  is 
a  complex  cyanide.  Several  different  ones  may  form,  depending  upon 
conditions,  a  typical  one  having  the  formula  KCu(NC)2.  This  com- 
pound ionizes  as  shown  in  the  equation 

KCu(NC)2^=>:K+  +  Cu(NC)g- 

The  copper  has  become  a  part  of  a  complex  anion,  and  only  to  an 
extremely  limited  extent  does  it  produce  simple  copper  ions  by  a 
secondary  ionization : 


FIG.  148 


Electric  cells.  An  electric  cell  is  a  device 
for  conyerting  chemical  energy  directly  into 
electrical  energy.  A  great  many  different 
chemical  reactions  can  be  arranged  in  such  a 
way  as  to  accomplish  this  result,  and  the  com- 
bination known  as  the  Daniell  cell  will  serve 
as  an  illustration  of  the  most  familiar  types  of 
cells.  In  this  combination  two  plates,  one  of 
copper  and  the  other  of  zinc,  each  fashioned 
so  as  to  have  a  large  surface,  are  arranged  in 
a  glass  jar,  as  shown  in  Fig.  148.  The  electrolyte  in  contact  with  the 
zinc  plate  is  zinc  sulfate,  while  that  in  contact  with  the  copper  plate 
is  copper  sulfate. 

The  action  of  the  cell  can  be  explained  as  follows :  The  zinc  atoms  have  a 
tendency  to  give  up  to  the  zinc  plate  A  two  electrons  each,  and  to  pass  into  solu- 
tion as  zinc  ions,  the  force  urging  this  change  being  designated  as  solution  ten- 
sion. But  since  the  zinc  ions  are  positively  charged,  and  their  formation  leaves 
the  zinc  plate  negatively  charged,  the  accumulation  of  these  charges  soon  pro- 
duces an  equilibrium  by  the  attraction  of  the  zinc  plate  for  the  positive  ions. 
Copper  ions,  on  the  other  hand,  tend  to  leave  the  solution  because  of  their  osmotic 
pressure,  and  to  deposit  as  metallic  atoms  upon  the  copper  plate  B,  each  copper 
ion  recovering  two  electrons  from  the  copper  plate.  Since  this  process  results  in 
charging  the  copper  plate  positively,  the  accumulated  charge  soon  produces  an 


486  GENEKAL  CHEMISTRY 

equilibrium  by  repelling  the  positive  copper  ions.  If  now  the  two  plates  are' 
joined  by  a  wire,  the  excess  electrons  on  the  zinc  plate  flow  through  the  wire  to 
make  up  the  deficiency  upon  the  copper  plate.  This  prevents  an  accumulated 
charge  on  either  plate  and  results  in  a  current  through  the  wire.  The  chemical 
action  taking  place  is  represented  by  the  equation 

Zn  +  CuSO4  =  Cu  +  ZnSO4  +  50,100  cal. 

in  which  nearly  all  the  heat  is  transformed  into  electrical  energy.  The  reaction 
ceases  when  the  wire  connection  is  broken. 

The  order  of  the  metals  in  the  electromotive  series  (p.  158)  is  the  order  of  in- 
tensity with  which  the  metals  tend  to  pass  into  ionic  form.  Any  two  metals  in  a 
suitable  electrolyte  will  constitute  a  cell  in  which  the  metal  highest  in  the  series 
is  the  negative  pole  and  the  lower  one  the  positive.  As  a  rule,  only  a  part  of 
the  chemical  energy  is  converted  into  electrical  energy,  the  remainder  being 
transformed  into  heat. 

MERCURY 

History  and  occurrence.  The  element  mercury,  or  quicksilver,  as  it 
is  usually  called,  was  known  considerably  before  the  Christian  era 
and  played  an  important  part  in  the  alchemy  of  the  Middle  Ages.  It 
is  found  in  a  number  of  localities,  usually  in  the  form  of  a  red  sulfide 
called  cinnabar  and  occasionally  as  drops  of  native  metal.  As  a  rule, 
the  ores  are  not  very  rich,  in  many  cases  carrying  less  than  1  per  cent  of 
mercury.  The  countries  which  produce  the  most  mercury  (in  the  order 
of  their  present  production)  are  Spain,  Italy,  Austria,  and  California. 

Metallurgy.  The  metallurgy  of  mercury  is  extremely  simple,  owing 
to  the  ready  decomposition  of  mercury  compounds  and  the  volatility 
of  the  metal.  It  is  only  necessary  to  roast  the  sulfide  in  a  current  of 
air  or  with  the  addition  of  calcium  oxide  : 


4  HgS  +  4  CaO  =  3  CaS  +  CaSO4  +  4  Hg 

The  resulting  vapor  of  mercury  is  easily  condensed  to  a  liquid. 

Purification  of  mercury.  Solid  materials  mixed  with  mercury  are 
removed  by  filtration  through  soft  leather.  The  impurities  remaining 
are  chiefly  other  metals  held  in  solution.  With  few  exceptions  these 
are  more  easily  oxidized  than  mercury  and  may  be  removed  by  digest- 
ing the  mercury  with  a  solution  of  an  oxidizing  agent,  such  as  nitric 
acid  or  ferric  chloride. 

Laboratory  purification.  In  the  laboratory  it  is  often  necessary  to  prepare  pure 
mercury  for  various  purposes,  and  this  may  be  accomplished  conveniently  by  the 
apparatus  shown  in  Fig.  149.  A  long  glass  tube  A,  drawn  out  to  an  S-shaped 
trap  B  at  the  lower  end,  is  filled  with  enough  mercury  to  close  the  trap  and  is  then 


COPPER;  MERCURY;  SILVER 


48T 


00 


filled  up  with  a  solution  of  ferric  chloride.  The  end  of  the  funnel  C  is  drawn  to 
a  fine  tip  through  which  the  mercury  streams,  falling  through  the  ferric  chloride 
solution  in  a  fine  spray  and  overflowing  at  the  bottom  into  a  receiving  vessel  D. 
For  very  refined  purposes  it  is  better  to  distill  the  mercury  under  diminished 
pressure,  though  Hulett  has  shown  that  the 
presence  of  a  little  oxygen  is  preferable  to  a 
complete  vacuum,  insuring  the  oxidation  of 
the  metallic  impurities  and  preventing  their 
volatilization  along  with  the  mercury. 

Properties.  Pure  mercury  is  a  silvery 
liquid  at  ordinary  temperatures,  and  to 
this  fact  it  owes  the  name  quicksilver  (the 
word  quick  meaning  "  live,"  or  "  mov- 
ing"). It  solidifies  at  -38.8°,  boils  at 
356.7°,  and  at  15°  has  a  density  of  13.56. 
Its  various  properties  give  it  great  value 
in  scientific  experimentation.  It  is  a 
convenient  liquid  over  which  to  collect 
gases  that  are  soluble  in  water.  It  has  a 
moderately  large  coefficient  of  expansion 
with  temperature  changes,  and  this  fact, 
together  with  its  low  freezing  point  and 
fairly  high  boiling  point,  renders  it  suit-  . 

able  for  use  in  the  construction  of  ther- 
mometers.   Its   density  and   low   vapor 

pressure  at  ordinary  temperatures  make  it  a  convenient  liquid  for 
barometers,  since  a  column  less  than  one  meter  in  height  will  balance 
the  atmospheric  pressure,  with  all  its  range  of  variation.  It  is  a  mod- 
erately good  conductor  of  electricity  and  therefore  affords  a  convenient 
means  of  joining  conducting  wires  in  cases  where  rigid  joints  would 
be  inconvenient. 

It  forms  alloys  (called  amalgams)  with  practically  all  the  metals. 
This  property  leads  to  the  largest  industrial  use  of  the  metal,  namely, 
in  the  extraction  of  gold  and  silver  from  their  ores.  When  little  of 
the  other  metal  is  present,  the  amalgams  are  liquid  but  not  so  mobile 
as  pure  mercury  and  inclined  to  be  stringy.  Some  of  these  amalgams 
have  industrial  uses.  Sodium  amalgam  is  an  efficient  reducing  agent. 

Chemical  conduct.  Mercury  is  a  metal  of  rather  feeble  chemical 
activity.  It  stands  low  in  the  electromotive  series  and  is  thrown  out 
of  combination  by  most  other  metals.  Nearly  all  of  its  compounds 


FIG.  149 


488  GENERAL  CHEMISTEY 

dissociate  at  high  temperatures,  yielding  free  mercury.  Heated  in  the 
air  at  temperatures  below  its  boiling  point,  it  slowly  combines  with 
oxygen  to  form  the  red  oxide,  but  this  is  easily  decomposed  at  higher 
temperatures,  the  reaction  being  reversible  : 


Under  ordinary  conditions  it  will  not  displace  hydrogen  from  acids, 
but  oxidizing  acids  attack  it,  forming  the  corresponding  salts.  Like 
copper,  its  affinity  for  sulfur  and  the  halogen  elements  is  stronger 
than  for  oxygen. 

Salts  of  mercury.  Like  copper,  mercury  forms  two  series  of  com- 
pounds —  the  mercurous  and  the  mercuric.  The  general  stability  of 
the  two  series  is  much  more  nearly  equal  than  is  the  case  with  copper 
salts,  and  salts  of  the  oxygen  acids,  as  well  as  of  practically  all  other 
acids,  are  known  in  each  series.  The  salts  of  mercury  are  remarkable 
for  the  fact  that,  compared  with  other  salts,  they  are  very  little  ionized 
in  solution.  For  example,  a  normal  solution  of  mercuric  chloride  at 
ordinary  temperatures  is  ionized  to  an  extent  of  less  than  0.01  per  cent. 
The  salts  of  'mercury  are  also  much  more  generally  soluble  in  organic 
solvents,  such  as  alcohol  and  ether,  than  is  usually  the  case  with 
metallic  salts.  Both  the  metal  and  its  salts  are  poisonous. 

Mercurous  salts.  Mercurous  salts  are  obtained  in  either  of  two 
general  ways  : 

1.  By  precipitation.  The  insoluble  salts  can  be  obtained  by  double 
decomposition  with  soluble  mercurous  salts,  such  as  the  nitrate  HgNOg: 

HgNO3  +  NaBr  =  HgBr  +  NaNO3 

2.  By  reduction  of  a  mercuric  salt.  This  is  most  conveniently  effected 
by  the  use  of  mercury  as  a  reducing  agent.    Thus,  mercuric  chloride, 
when  heated  with  mercury,  yields  mercurous  chloride  : 


When  mercury  is  dissolved  in  cold,,  dilute  oxidizing  acids,  the  mer- 
curous salt  is  obtained  if  mercury  is  in  excess,  for  any  mercuric  salt 
which  forms  is  reduced  by  the  mercury. 

The  molecular  weight  of  mercurous  compounds.  There  has  been  much  discus- 
sion as  to  the  true  molecular  weights  of  mercurous  salts,  as  well  as  of  those  of 
cuprous  salts  and  ferric  salts,  and  these  compounds  are  sometimes  given  double 
formulas,  such  as  Hg2Cl2,  Fe2Cl6.  The  reason  for  this  is  in  part  the  influence  of 
an  old  theory,  long  since  abandoned,  and  in  part  the  evidence  drawn  from  actual 


COPPER;  MERCURY;  SILVER  489 

determinations  of  molecular  weights.  The  latter  evidence  is  not  always,  easy  of 
interpretation,  for  the  vapor  density  varies  with  the  temperature,  and  the  measure- 
ments are  complicated  because  of  dissociation,  such  as  occurs  in  the  case  of 
mercurous  chloride:  2  HgCl  =  HgCl2  +  Hg 

Inferences  drawn  from  solution  measurements  (boiling  and  freezing  points)  are 
also  unreliable,  for  there  is  evidence  that  in  solution  some  molecules  may  poly- 
merize, while  others  ionize.  In  the  case  of  many  salts  which  are  not  volatile  and 
are  insoluble,  we  have  no  evidence  at  all.  It  seems  best,  therefore,  to  retain  the 
simple  formulas  in  the  case  of  all  .salts,  especially  since  no  method  gives  us 
information  as  to  the  true  molecular  weight  in  the  solid  state. 

Mercurous  oxide  (Hg20) ;  hydroxide  (HgOH) ;  sulfide  (Hg2S).  All  of 
these  insoluble  compounds  are  apparently  formed  by  the  usual  methods 
of  preparation,  but  they  are  very  unstable  and  quickly  decompose  into 
the  more  stable  mercuric  compounds,  especially  under  the  influence  of 
sunlight.  The  preparation  and  decomposition  are  illustrated  in  the 
equation  for  the  sulfide: 

HgJSO.  +  H2S  =  H2S04  +  Hg2S  — *•  HgS  +  Hg 

Mercurous  halides.  The  mercurous  halogen  compounds  are  com- 
paratively stable  and  well  characterized,  the  iodide  being  the  least  so. 
They  are  insoluble  in  water.  The  chloride  (HgCl),  known  as  calomel, 
is  prepared  by  subliming  mercuric  chloride  with  mercury : 

HgCl2  +  Hg  =  2HgCl 

It  is  also  prepared  by  subliming  mercurous  sulfate  with  common 
salt.  It  is  a  white,  crystalline  body,  easily  volatile,  and  has  important 
uses  as  a  drug.  In  strong  sunlight  the  reaction  just  given  is  to  some 
extent  reversed,  the  preparation  darkening  in  consequence  of  the  sepa- 
ration of  mercury.  Since  the  mercuric  chloride  formed  at  the  same 
time  is  exceedingly  poisonous,  it  is  necessary  to  preserve  calomel  in 
dark  bottles.  The  bromide  (HgBr)  and  the  iodide  (Hgl)  have  similar 
properties.  They  decompose  with  ease,  the  iodide  ranging  in  color 
from  yellow  to  green,  apparently  through  decomposition. 

Mercurous  nitrate  (HgN03).  This  salt  is  formed  when  cold,  dilute 
nitric  acid  acts  upon  mercury.  It  is  quite  soluble,  forms  monoclinic 
needles  of  the  hydrate  HgNO3  •  2  H2O,  and  undergoes  hydrolysis  in 
dilute  solution,  forming  a  basic  salt. 

Mercurous  sulfate  (Hg2SOJ.  Mercurous  sulfate  is  formed  in  a  similar 
way  by  the  action  of  sulf  uric  acid  upon  mercury.  A  very  pure  prepa- 
ration may  be  made  by  placing  a  suitable  quantity  of  mercury  in  a 


490 


GENERAL  CHEMISTRY 


beaker  and  pouring  over  it  dilute  sulfuric  acid.  A  wire  is  dipped 
into  the  mercury,  which  is  made  to  serve  as  anode,  while  a  small  piece 
of  platinum  foil  dipped  into  the  acid  serves  as  the  cathode.  When 
a  suitable  current  is  passed  through  the  cell  so  formed,  the  mercury 
dissolves  in  the  acid,  forming  mercurous  sulfate.  The  mercury  is 
gently  stirred  during  solution,  to  prevent  the  formation  of  the  mercuric 
salt.  The  sulfate  is  somewhat  hydrolyzed  by  water. 

Standard  cells.  The  chief  use  of  mercurous  sulfate  is  in  the  construction  of 
standard  cells.  These  are  small  cells  which  yield  a  very  constant  electromotive 

force,  against  which  other  cells  may  be  stand- 
ardized. The  arrangement  most  often  used 
is  represented  in  Fig.  150  and  is  called  the 
Weston  cell.  Two  small  glass  test  tubes, 
through  the  bottom  of  each  of  which  is  sealed 
a  platinum  wire,  are  connected  by  a  glass 
tube  so  as  to  form  an  H-shaped  vessel.  Mer- 
cury in  contact  with  a  paste  of  mercurous 
sulfate  and  water  is  placed  in  the  one  tube 
(.4),  while  cadmium  (in  the  form  of  amal- 
I  ~|  gam),  in  contact  with  a  saturated  solution  of 

•pIG   jgQ  cadmium  sulfate,  is  placed  in  the  other  (£), 

the  connecting  tube  being  also  filled  with 

the  latter  solution.  Crystals  of  cadmium  sulfate  are  placed  in  each  tube  to  in- 
sure the  saturation  of  the  solution.  Such  a  cell  has  an  electromotive  force  of 
1.0186  volts  at  20°. 

Mercuric  salts.  Mercuric  salts  are  usually  prepared  by  oxidation  of 
mercurous  salts  or  by  the  solution  of  mercuric  oxide  or  mercury  in 
the  appropriate  acid.  As  a  rule,  they  are  more  soluble  than  the  cor- 
responding mercurous  salts  and  are  more  extensively  hydrolyzed  in 
solution,  yielding  a  great  variety  of  basic  salts,  which  are  usually  of 
some  shade  of  yellow.  They  also  yield  many  complex  compounds. 

Mercuric  oxide  (HgO).  Mercuric  oxide  is  prepared  as  a  bright  red, 
crystalline  powder  by  the  careful  heating  of  mercuric  nitrate.  It  is 
also  obtained  as  a  yellow  precipitate  when  a  cold  solution  of  a  soluble 
base  is  poured  into  a  solution  of  a  mercuric  salt.  The  hydroxide, 
which  would  be  expected  to  form  in  the  reaction,  spontaneously 
decomposes  into  oxide: 

Hg(N03)2  +  2  KOH  =  2  KNO3  +  Hg(OH)2  — >-  HgO  +  H2O 

The  yellow  oxide  changes  into  the  red  at  a  higher  temperature,  but  it 
is  not  entirely  certain  whether  these  two  are  distinct  forms  or  whether 
they  owe  their  different  colors  to  differences  in  their  fineness  of  division. 


COPPER.;  MERCURY;  SILVER  491 

When  the  oxide  is  heated,  it  dissociates  into  mercury  and  oxygen  — 
a  reaction  which  led  Priestley  to  the  discovery  of  oxygen. 

Mercuric  sulfide  (cinnabar)  (HgS).  Mercuric  sulfide  occurs  in  nature 
as  the  red  mineral  cinnabar,  and  is  the  most  important  natural  com- 
pound of  mercury.  The  compound  can  be  obtained  by  precipitation, 
as  shown  in  the  following  equation : 

Hg(N08)2  +  H2S  =  HgS  +  2  HNO, 

When  it  is  so  prepared,  it  is  a  black,  amorphous  substance,  very 
highly  insoluble  in  water  and  in  acids.  When  mercury  and  sulfur  are 
intimately  ground  together  (forming  the  black,  amorphous  sulfide) 
and  the  product  is  warmed  with  a  solution  of  potassium  sulfide  (in 
which  mercuric  sulfide  is  somewhat  soluble),  the  bright  red  variety 
known  as  vermilion  is  obtained.  It  is  a  valuable  pigment. 

Mercuric  chloride  (corrosive  sublimate)  (PgCl2).  Mercuric  chloride 
can  be  obtained  by  the  usual  methods  as  a  white,  crystalline  solid 
moderately  soluble  in  water.  In  the  industries  it  is  made  by  heating 
a  mixture  of  ordinary  salt  and  mercuric  sulfate,  and  condensing  the 
volatile  chloride.  It  is  soluble  in  alcohol  and  in  ether,  as  well  as 
in  water,  and  is  extensively  used  in  surgery  as  an  antiseptic.  It  is 
extremely  poisonous.  It  combines  directly  with  a  great  variety  of 
substances,  among  others  with  albumin.  This  latter  fact  leads  to  the 
use  of  white  of  egg  as  an  antidote  for  the  compound.  When  treated 
with  a  reducing  agent,  it  is  easily  reduced  to  mercurous  chloride  and 
often  to  free  mercury.  The  reactions  with  stannous  chloride,  for 
example,  are  shown  in  the  following  equations : 

2  HgCl0  +  SnCl2  =  2  HgCl  +  SnCl4 
2  HgCl  +  SnCl2  =  2  Hg  +  SnCl4 

Mercuric  iodide  (HgI2).  This  salt  is  interesting  as  occurring  in  two 
very  different  forms.  At  ordinary  temperatures  the  stable  form  is 
bright  scarlet  in  color.  Above  126.5°  this  changes  to  a  yellow  modi- 
fication which  melts  at  223°,  and  when  cooled  below  126.5°  it  changes 
again  into  the  red  form.  It  is  very  sparingly  soluble  in  water,  but 
dissolves  readily  in  solutions  of  potassium  iodide  or  in  potassium 
hydroxide,  forming  the  complex  compound  K2HgI4.  This  solution  in 
potassium  hydroxide  is  known  as  Nessler's  reagent,  and  it  is  used  in 
detecting  the  presence  of  small  traces  of  ammonia,  especially  in  potable 
waters.  The  reactions  will  be  described  in  a  later  paragraph. 


492  GENERAL  CHEMISTRY 

Mercuric  cyanide  (Hg(NC)2).  This  compound  is  prepared  by  dissolv- 
ing mercuric  oxide  in  hydrocyanic  acid.  It  is  a  white,  well-crystallized 
salt  and  is  remarkable  among  the  cyanides  of  the  heavier  metals  in 
being  soluble  both  in  water  and  in  alcohol.  When  heated  it  decomposes 
into  mercury  and  cyanogen  : 


Other  mercuric  salts.  A  few  other  mercuric  salts  should  be  mentioned 
very  briefly. 

Mercuric  nitrate  (Hg(NO8)2)  crystallizes  from  dilute  nitric  acid  in 
the  form  of  a  hydrate  of  the  composition  Hg(NOg)2  •  8  H2O.  In  water 
it  undergoes  hydrolysis,  with  the  formation  of  a  number  of  basic  salts. 

Mercuric  sulfate  (HgSO4)  is  known  both  as  the  anhydrous  salt  and 
as  the  monohydrate  HgSO4  •  H2O.  Both  salts  are  white,  crystalline 
solids,  which  are  hydrolyzed  in  the  presence  of  much  water,  forming 
a  yellow  basic  salt  of  the  formula  HgSO4  •  2  HgO,  known  as  turpeth. 

Mercuric  fulminate  (Hg(ONC)2)  is  prepared  by  the  action  of  nitric 
acid  upon  mercury  in  the  presence  of  alcohol.  It  is  extremely  explosive 
and  is  used  in  the  manufacture  of  percussion  cartridges. 

Ammonia  compounds  of  mercury.  Among  the  most  numerous  com- 
plex compounds  of  mercury  are  those  which  its  salts  form  with  ammonia. 
More  than  a  hundred  of  these  compounds  have  been  described,  and 
these  have  been  extensively  studied  by  Franklin.  They  can  be  under- 
stood most  easily  by  remembering  that  liquid  ammonia  and  water, 
as  solvents,  have  many  qualities  in  common,  and  that  their  reactions 
with  salts  are  closely  analogous.  These  compounds  may  then  be 
grouped  into  three  classes: 

1.  Analogues  of  hydrates.  Just  as  a  salt  may  take  up  water  to  form 
a  hydrate,  so  it  may  combine  with  ammonia  to  form  an  ammoniate  : 

CuCl2  +  2  NH3  =  CuCl2  •  2  NH3 
HgCl2  +  2  NH3  =  HgCl2  -  2  NH8 

The  ammoniate  of  calcium  chloride   (CaCl2  •  8  NH8)  has  been  men- 
tioned in  earlier  pages  (p.  172). 

2.  Analogues  of  basic  salts.  The  hydrolysis  of  salts  has  been  referred 
to  in  a  number  of  places  and  is  well  illustrated  in  the  case  of  bismuth 
chloride,  in  which  the  following  stages  occur  : 

/Cl  /OH  /OH  ,0 

Bi^Cl     -  —  *•     Bi-Cl       -  >-     Bi-OH     -  >-    Bi^      +  H2O 
XC1  XC1  XC1  XC1 


COPPER;  MERCURY;  SILVER  493 

With  mercuric  salts  there  is  a  corresponding  action  which  Franklin 
has  aptly  named  ammonolysis.    With  the  chloride  the  stages  are 


— >  HgNH  +  NH8 

3.  Mixed  types.  When  mercury  salts  are  treated  with  aqueous  am- 
monia, which  may  be  regarded  as  a  mixed  solvent  consisting  of  water 
and  ammonia,  both  hydrolysis  and  ammonolysis  may  occur.  In  the 
case  of  the  chloride  this  is  represented  by  the  equation 


-  O-H  „  OH 

NH 


Hg/          +2HC1 


Reactions  which  can  be  referred  to  these  types  occur  with  most  mercuric  salts. 
With  mercurous  salts  there  is  at  the  same  time  a  decomposition  which  results  in 
the  formation  of  a  mercuric  derivative  and  free  mercury,  as  is  illustrated  in  the 

case  of  calomel : 

Pi 
2  HgCl  +  XH3  -  Hg  <  ^H   +  HC1  +  Hg 

(Cl     \ 
Hg  <  -^jj  )  formed  in  this  reaction  is  a  white 

compound,  but  it  appears  to  be  jet-black,  owing  to  the  finely  divided  mercury 
precipitated  at  the  same  time.  Many  of  these  compounds  were*  known  long  be- 
fore their  nature  was  understood,  and  they  have  various  accidental  names.  A 
few  well-known  ones  are  the  following : 

HgCl2  •  2  NH3  :  fusible  white  precipitate  (class  1). 
NH2  —  Hg  —  Cl :  infusible  white  precipitate  (class  2). 
HO  -  Hg  -  NH  -  Hg  -  OH  :  Millon's  base  (class  3). 
HO  -  Hg  -  NH  -  Hg  -  I :  Nessler's  precipitate. 

This  last  compound  is  the  iodide  of  Millon's  base  and  is  formed  when  mercuric 
chloride,  potassium  iodide,  and  ammonia  are  brought  together  in  alkaline  solu- 
tion. It  is  a  dark,  reddish-brown  precipitate,  and  even  minute  traces  of  ammonia 
will  give  a  yellow  coloration  with  these  reagents  (Xessler's  solution). 

Other  reactions  of  ammonia  with  salts.  In  addition  to  the  reactions  just  de- 
scribed, it  will  be  recalled  that  solutions  of  ammonia  may  act  upon  metallic  salts 
in  either  of  two  ways  : 

1.  By  the  precipitation  of  an  insoluble  hydroxide.  This  is  the  most  familiar  case  and 
is  illustrated  by  the  action  of  aqua  ammonia  upon  ferric  chloride : 

FeCl3  +  3  NH4OH  =  Fe(OH)8  +  3  NH4C1 

2.  By  the  formation  of  a  soluble  complex  salt.  This  action  has  been  described  in  con- 
nection with  the  salts  of  cobalt  and  copper.    With  copper  sulfate  the  equation  is 
as  follows : 

CuSO4  +  4  NH3  =  Cu(NH3)4SO4 


494  GENERAL  CHEMISTRY 

SILVER 

History  and  occurrence.  Silver,  the  argentum  of  the  Romans,  has 
been  known  from  the  earliest  times  and,  together  with  gold,  has 
always  ranked  as  a  precious  metal.  It  is  frequently  found  native  in 
the  form  of  flakes  or  wire  imbedded  in  primitive  rock,  and  occasion- 
ally in  large  masses.  In  the  combined  state  it  occurs  in  many  minerals, 
those  of  most  importance  to  the  metallurgist  being  the  following :  cerar- 
gyrite  (horn  silver)  (AgCl);  argentite  (Ag2S);  proustite  (Ag3AsS3); 
pyrargyrite  (AggSbS3) ;  stephanite  (5  Ag2S  •  Sb2S3).  It  is  also  found  in 
small  quantities  in  practically  all  copper  and  lead  ores,  a  considerable 
quantity  of  the  silver  now  produced  coming  from  this  source. 

Metallurgy.  In  a  text  of  this  scope  it  is  possible  to  explain  only  the 
most  general  principles  of  the  metallurgy  of  silver,  since  the  details 
are  very  complicated  and  subject  to  frequent  change.  The  methods 
employed  may  be  classified  under  three  heads : 

1.  Smelting.  In  furnace  smelting  the  ores  are  mixed  with  lead  ores 
(if  they  are  not  already  rich  enough  in  lead),  and  the  two  metals  are 
obtained  as  an  alloy.    The  separation  of  the  silver  is  described  in 
connection  with  lead  (p.  505). 

2.  Amalgamation.  When  silver  occurs  native,  or  in  forms  which  can 
be  readily  converted  into  metallic  silver  by  suitable  reagents,  the  ores, 
together  with  the  reagents,  are  thoroughly  mixed  with  mercury,  which 
dissolves  the  silver  and  gold.    The  resulting  amalgam  is  then  collected 
and  the  mercury  distilled  off,  leaving  the  impure  silver. 

3.  Hydrometallurgy.  In  this  process  the  silver  is  dissolved  from  the 
finely  crushed  ore  by  a  solution  of  some  suitable  reagent.    Sometimes 
sulfide  ores  are  roasted  until  the  silver  is  converted  into  sulfate,  which 
is  itself  soluble  in  water.    In  other  cases  solutions  of  sodium  cyanide, 
potassium-mercurous  cyanide  (KHg(NC)2),  or  sodium  thiosulfate  are 
employed,  all  of  which  dissolve  silver,  as  well  as  some  silver  compounds! 
From  the  solution  of  silver  so  obtained  the  metal  is  precipitated  by  a 
suitable  reagent,  such  as  copper,  zinc,  or  sodium  sulfide. 

Refining  of  silver.  There  are  a  number  of  methods  by  which  silver 
is  refined. 

1.  Cupellation  and  parting  with  sulfuric  acid.  In  this  process  the  metal 
is  heated  on  an  open  hearth  in  a  strong  current  of  air.  The  various 
metallic  impurities  (excepting  gold)  are  in  this  way  largely  converted 
into  oxides  and  swept  off  as  dross,  leaving  the  silver  alloyed  with 


COPPEK;  MERCURY;  SILVER  495 

small  percentages  of  gold,  copper,  and  iron.    It  is  then  cast  into  ingots 
known  as  dore  bars,  since  they  contain  gold. 

In  order  to  recover  the  gold,  the  alloy  is  treated  with  hot  concen- 
trated sulfuric  acid,  which  converts  all  the  metals,  except  the  gold, 
into  sulf  ates.  When  water  is  added  to  the  resulting  mixture,  the 
sulfates  of  copper,  silver,  and  iron  pass  into  solution,  while  the  gold, 
together  with  the  lead  sulfate  and  any  unattacked  substances,  settles 
as  a  mud  from  which  the  gold  is  subsequently  recovered.  The  silver 
is  separated  from  the  solution  of  the  sulfates  by  suspending  in  the 
latter  clean  copper  plates,  the  copper  displacing  the  silver,  which  is 
deposited  in  crystalline  form  : 

4  +  Cu  =  CuS04  +  2  Ag 


The  copper  sulfate  obtained  as  a  by-product  in  this  process  furnishes 
much  of  the  blue  vitriol  of  commerce. 

2.  Electrolytic  refining.  Electrolysis  of  the  impure  silver  is  now  car- 
ried out  extensively,  the  process  being  conducted  in  a  way  very  similar 
to  the  electrolysis  of  copper.  The  electrolyte  used  is  a  solution  of 
silver  nitrate  in  nitric  acid.  The  silver  is  deposited  as  crystals,  which 
are  mechanically  brushed  off  the  cathode,  collected,  and  melted 
into  bars. 

The  United  States  produces  about  one  third  of  the  world's  out- 
put of  silver,  and  America,  including  Mexico  and  Canada,  about 
70  per  cent. 

Properties.  Silver  is  a  brilliant  white  metal  which  melts  at  960°,  boils 
at  1955°,  and  has  a  density  of  10.5.  It  is  very  ductile  and  malleable 
and  has  the  greatest  electrical  conductivity  of  all  the  metals.  It  is 
intermediate  in  hardness  between  gold  and  copper,  and  in  thin  foil 
transmits  blue  light.  It  alloys  with  many  other  metals  and  dissolves 
readily  in  mercury,  forming  an  amalgam.  When  melted  it  dissolves 
notable  quantities  of  oxygen,  giving  it  up  again  during  solidification, 
with  a  characteristic  sputtering  of  the  metal.  One  gram  of  silver  gives 
up  about  1  cc.  of  oxygen.  It  crystallizes  in  octahedra,  but  much  of 
the  native  silver  is  amorphous.  It  can  be  obtained  in  colloidal  sus- 
pension by  sparging  silver  wires  under  water  (p.  132)  or  by  chemical 
reduction.  It  is  extensively  used  for  household  ornaments  and  utensils, 
for  coinage  (p.  480),  in  the  manufacture  of  mirrors,  and  as  halogen 
salts  in  photography.  Much  of  the  output  is  exported  each  year  to 
the  Far  East. 


496  GENERAL  CHEMISTRY 

Chemical  conduct.  Silver  is  one  of  the  inactive  elements.  It  does 
not  combine  directly  with  oxygen,  hydrogen,  carbon,  nitrogen^silicon, 
and  other  elements,  nor  is  it  acted  upon  by  dilute  acids  .or  fused 
alkalies.  The  halogen  elements  attack  it  slowly  and  only  superficially, 
since  insoluble  halides  are  formed  which  prevent  further  action.  Ozone 
attacks  it,  forming  a  black  peroxide.  Sulfur,  as  well  as  most  sulfur 
compounds,  blackens  it,  owing  to  the  formation  of  silver  sulfide  (oxi- 
dized silver).  The  oxidizing  acids  dissolve  it,  forming  the  correspond- 
ing salts,  such  as  the  nitrate  (AgNOg)  and  the  sulfate  (Ag2SO4). 
Potassium  cyanide,  in  the  presence  of  air  and  water,  dissolves  it 
according  to  the  following  equation: 

8  KNC  +  O2  +  2  H2O  +  4  Ag  =  4  KAg(NC)2  +  4  KOH 

Compounds  of  silver.  Silver  forms  only  one  well-defined  series  of 
salts,  and  in  these  it  is  univalent.  These  salts  are  usually  colorless  or 
light  yellow,  and  are  little  hydrolyzed  in  solution,  yielding  few  basic 
salts.  They  are  readily  reduced  to  metallic  silver,,;  and  consequently 
are  mild  oxidizing  agents.  In  addition  to  these  salts  there  are  a  few 
compounds  in  which  the  valence  is  open  to  question,  and  there  are 
many  complex  salts.  The  salts  of  silver  are  characterized  by  the  fact 
that  very  few  of  them  form  hydrates. 

Silver  oxide  (Ag20).  Silver  oxide  is  thrown  down  as  a  dark  brown, 
amorphous  precipitate  when  a  soluble  hydroxide  is  added  to  a  solution 
of  a  silver  salt,  the  hydroxide  at  first  formed  spontaneously  decom- 
posing into  water  and  the  oxide.  It  is  sufficiently  soluble  in  water  to 
turn  red  litmus  blue,  evidently  forming  some  hydroxide,  and  it  acts 
as  a  strong,  highly  ionized  base,  yielding  salts  which  are  neutral  to 
litmus.  At  a  very  moderate  temperature  it  decomposes  into  silver 
and  oxygen,  the  reaction  being  rapid  at  250°. 

Other  oxides.  A  suboxide  (Ag4O)  has  been  described  as  being  formed  when 
certain  silver  compounds  are  carefully  reduced.  There  is  also  evidence  for  the 
existence  of  a  series  of  compounds  derived  from  this  oxide,  of  which  the  fluoride 
(Ag2F)  is  crystalline,  while  the  others  (Ag2Cl,  Ag2Br,  Ag2I,  and  Ag4S)  are  less 
well  defined.  A  peroxide  (Ag2O2)  is  formed  by  the  action  of  ozone  upon  silver. 

Silver  halides.  Silver  fluoride  (Ag2F2)  is  prepared  by  dissolving 
silver  oxide  in  hydrofluoric  acid.  It  is  a  very  soluble  salt,  crystallizing 
as  the  hydrate  Ag2F2  •  2  H2O  or  Ag2F2  •  4  H2O.  The  other  halides,  on 
the  contrary,  are  practically  insoluble,  the  chloride  being  the  most 
soluble  of  the  three  and  the  iodide  the  least  so.  They  are  prepared 


COPPER;  MERCURY;  SILVER  497 

by  treating  a  soluble  silver  salt  with  either  a  soluble  halide  salt  or 
the  free  acid  :  ^^  +  ^  =  ^  +  ^^ 

They  form  amorphous,  curdy  precipitates,  the  chloride  being  white, 
the  bromide  very  pale  yellow,  the  iodide  decidedly  yellow.  All  three 
are  also  known  in  crystalline  condition.  The  formation  of  these  pre- 
cipitates is  used  as  a  test,  on  the  one  hand  for  silver  ions  and  on  the 
other  for  the  halogen  ions.  The  halides  are  soluble  in  a  number  of 
normal  salts,  forming  a  series  of  complexes.  This  property  is  of  impor- 
tance in  metallurgy  and  in  photography.  A  few  of  these  reactions  are 
shown  in  the  following  equations,  silver  chloride  being  used  as  an 
example  : 


RC1 

AgCl  +  2  Na2S203  =  Na3Ag(S2O3)2  '+  NaCl 

The  chloride  is  also  soluble  in  ammonium  hydroxide,  forming  the 
complex  salt  Ag(NHg)2Cl,  which  can  be  obtained  in  the  form  of 
shining,  colorless  needles.  With  gaseous  ammonia  the  chloride  and 
iodide  form  the  addition  compounds  2  AgCl  •  3  NHg,  AgCl  •  3  NHg, 
and  AgI-2NH,. 

All  three  of  the  insoluble  silver  halides  are  sensitive  to  light,  dark- 
ening in  color  as  the  result  of  a  change  in  which  a  portion  of  the 
halogen  element  is  liberated.  In  the  dark  this  reaction  is  reversed. 
While  the  nature  of  the  change  is  not  thoroughly  understood,  it  seems 
probable  that  it  is  represented  by  the  following  equation  : 

4AgCl+=±2Ag2Cl  +  Cl2 

It  is  upon  this  property  of  these  halides  that  the  art  of  photography 
is  based. 

Photography.  From  a  chemical  standpoint  the  processes  involved  in  photog- 
raphy may  be  described  under  two  heads  :  (1)  the  preparation  of  the  negative  ; 
(2)  the  preparation  of  the  print. 

1.  Preparation  of  the  negative.  The  plate  used  in  the  preparation  of  the  negative 
is  made  by  spreading  a  thin  layer  of  gelatin,  in  which  colloidal  silver  bromide  is 
suspended  (silver  iodide  is  sometimes  added  also),  over  a  glass  plate  or  celluloid 
film  and  allowing  it  to  dry.  When  the  plate  so  prepared  is  placed  in  a  camera 
and  the  image  of  some  object  is  focused  upon  it,  the  silver  salt  undergoes  a  change 
which  is  proportional  at  each  point  to  the  intensity  of  the  light  falling  upon  it. 
In  this  way  an  image  of  the  object  photographed  is  produced  upon  the  plate, 
which  is,  however,  invisible  and  is  therefore  called  latent.  It  can  be  made  visible 
by  the  process  of  developing. 


498  GENERAL   CHEMISTRY 

To  develop  the  image,  the  exposed  plate  is  immersed  in  a  solution  of  some  re- 
ducing agent,  called  the  developer.  While  the  developer  will  in  time  reduce  all  the 
silver  salt  present,  it  acts  much  more  rapidly  upon  that  which  has  been  exposed 
to  the  light.  The  plate  is  therefore  left  in  contact  with  the  developer  only  long 
enough  to  properly  bring  out  the  image.  The  resulting  metallic  silver  is  deposited 
in  the  form  of  a  black  film,  which  adheres  closely  to  the  plate. 

The  unaffected  silver  salt  is  then  removed  from  the  plate  by  immersing  it 
in  a  solution  of  sodium  thiosulfate.  After  the  silver  salt  has  been  dissolved,  the 
plate  is  washed  with  water  and  dried.  The  plate  so  prepared  is  called  the  nega- 
tive, because  it  is  a  picture  of  the  object  photographed,  with  the  lights  exactly 
reversed. 

2.  Preparation  of  the  print  The  print  is  made  from  paper  which  is  prepared  in 
the  same  way  as  the  negative  plate.  The  negative  is  placed  upon  this  paper  and 
exposed  to  the  light  in  such  a  way  that  the  light  must  pass  through  the  negative 
before  striking  the  paper.  If  the  paper  is  coated  with  silver  chloride,  a  visible 
image  is  produced,  in  which  case  a  developer  is  not  needed.  It  is  in  this  way  that 
proofs  are  made ;  in  order  to  make  them  permanent,  the  unchanged  silver  chloride 
must  be  dissolved  off  with  sodium  thiosulfate.  The  print  is  then  toned  by  dip- 
ping it  into  a  solution  of  gold  or  platinum  salts.  The  silver  on  the  print  passes 
into  solution,  while  the  gold  or  platinum  takes  its  place.  These  metals  give  a 
characteristic  color,  or  tone,  to  the  print,  the  gold  making  it  reddish  brown,  while 
the  platinum  gives  it  a  steel-gray  tone.  If  a  silver  bromide  paper  is  used  in  mak- 
ing the  print,  a  latent  image  is  produced  which  must  be  developed,  as  in  the  case 
of  the  negative  itself.  The  silver  bromide  is  much  more  sensitive  than  the  chlo- 
ride, so  that  the  printing  can  be  done  in  artificial  light.  Since  the  darkest  places 
on  the  negative  cut  off  the  most  light,  it  is  evident  that  the  lights  of  the  print 
will  be  the  reverse  of  those  of  the  negative,  and  will  therefore  correspond  to  those 
of  the  object  photographed. 

Silver  nitrate  (lunar  caustic)  (AgN03).  Silver  nitrate  is  prepared  by 
the  action  of  nitric  acid  upon  silver,  and  it  is  the  salt  of  silver  most 
extensively  used  in  the  laboratory.  It  is  extremely  soluble  in  water, 
100  g.  of  solvent  at  20°  dissolving  215  g.,  and  at  100°,  910  g.  It 
crystallizes  in  colorless  rhombic  plates  and  melts  undecomposed  at 
208.6°.  It  is  often  cast  into  sticks  for  use  as  a  caustic  in  surgery, 
its  use  depending  upon  the  fact  that  it  is  a  powerful  oxidizing  agent. 
It  produces  a  black  stain  on  the  skin,  owing  to  a  deposit  of  metallic 
silver.  A  similar  stain  is  produced  upon  any  other  oxidizable  material, 
such  as  cloth,  and  because  of  this  fact  silver  nitrate  is  used  in  the 
manufacture  of  indelible  inks.  It  dissolves  in  aqua  ammonia,  forming 
the  complex  Ag(NH3)2NOg. 

Other  soluble  salts.  Among  the  other  soluble  salts  of  silver  the 
most  important  are  the  sulfate  (Ag2SO4),  which  is  but  sparingly  solu- 
ble, and  the  acetate  (AgC2HgO2),  which  crystallizes  in  shining  needles. 
The  nitrite  (AgNO  )  is  soluble  in  hot  water. 


COPPER;  MERCURY;  SILVER 


499 


Insoluble  salts.  In  addition  to  the  halogen  compounds  already  de- 
scribed, many  of  the  salts  of  silver  are  insoluble  and  'have  charac- 
teristic colors,  or  appearances,  which  serve  to  identify  the  various 
anions  in  analysis.  Among  these  are  the  following: 

Silver  sulfide  (Ag2S)  is  found  in  nature  as  argentite  and,  as  prepared 
by  precipitation,  is  a  black,  amorphous  solid.  It  is  the  most  insoluble 
of  silver  compounds,  both  in  water  and  in  acids. 

Silver  cyanide  (AgNC)  forms  as  a  curdy  white  precipitate  when  a 
soluble  cyanide  is  added  to  a  silver  salt.  It  is  soluble  in  excess  of  the 
precipitant,  forming  a  complex  cyanide : 

AgNC  +  KNC  =  KAg(NC)2 

The  latter  plays  an  important  part  in  silver  plating. 
The  following  compounds  are  of  less  importance : 
Silver  carbonate  (Ag2COg)  :  a  pale  yellow  powder. 
Silver  chromate  (Ag2CrO4)  :  a  brick-red,  amorphous  solid. 
Silver  phosphate  (Ag3PO4)  :  a  clear,  yellow  solid. 
Silver  pyrophosphate  (Ag4P2O7)  :  a  granular,  white  powder. 

Electroplating.  Since  silver  has  a  pleasing  appearance  and  is  not  acted  upon 
by  water  or  air,  it  is  used  to  coat  various  articles  made  of  cheaper  metals.  Such 
articles  are  said  to  be  silver- 
plated.  The  process  by  which 
this  is  done  is  called  electro- 
plating. It  is  carried  on  as  fol- 
lows :  The  object  to  be  plated 
(a  spoon,  for  example)  is  at- 
tached to  a  wire  and  dipped 
into  a  solution  of  a  silver  salt. 
Electrical  connection  is  made 
in  such  a  way  that  the  article  to  be  plated  serves  as  the  cathode,  while  the  anode 
is  made  up  of  one  or  more  plates  of  silver  (Fig.  151,  A).  When  a  current  is 
passed  through  the  electrolyte,  silver  dissolves  from  the  anode  plate  and  deposits 
on  the  cathode  in  the  form  of  a  closely  adhering  layer.  By  making  the  proper 
change  in  the  electrolyte  and  the  anode  plate,  objects  may  be  plated  with  gold 
and  other  metals. 


FIG.  151 


CHAPTER  XXXVI 

TIN  AND  LEAD 

General.  The  elements  of  smaller  atomic  weight  in  Group  IV  of 
the  periodic  classification,  including  carbon,  silicon,  titanium,  and  zir- 
conium, are  acid-forming  in  character  and  have  already  been  described. 
The  elements  of  larger  atomic  weight,  of  which  tin  and  lead  are  the 
well-known  representatives,  are  essentially  metals.  Germanium  is  in- 
termediate in  character  and  is  so  rare  as  to  require  little  comment, 
while  thorium  is  a  metal  and  is  more  abundant.  These  four  elements 
are  not  so  closely  related  to  each  other  as  are  the  members  of  some 
other  families,  and  each  has  its  own  peculiarities.  The  more  abundant 
elements,  tin  and  lead,  will  be  described  first. 

TIN 

History  and  occurrence.  It  is  difficult  to  decide  just  when  tin  be- 
came known  to  the  ancients,  but  it  was  undoubtedly  at  a  very  early 
day.  In  ancient  times  it  was  frequently  confused  with  lead,  the  dis- 
tinction between  the  two  first  clearly  appearing  in  the  writings  of 
Pliny,  about  the  beginning  of  the  Christian  era.  It  is  found  chiefly  as 
the  oxide  SnO2,  called  cassiterite,  or  tinstone,  and  this  is  the  only  com- 
mercial ore.  It  has  long  been  produced  in  the  East  Indies  and  in 
Cornwall,  and  Bolivia  now  supplies  a  large  quantity.  Its  production 
in  the  United  States  is  insignificant. 

Metallurgy.  Since  tin  occurs  as  the  oxide  and  is  relatively  free 
from  other  metals,  its  metallurgy  is  very  simple  and  consists  in  reduc- 
ing the  ore  with  carbon.  In  some  cases  the  ore  is  first  roasted,  to  free 
it  from  sulfur  and  arsenic. 

Properties.  Tin  occurs  in  a  number  of  physical  modifications,  which 
pass  into  each  other  at  definite  transition  temperatures,  the  order  being 
as  follows:  (lg0)  (1610)  (2g20) 

Gray  tin    <    >    tetragonal    <    >   rhombic    <    >    liquid 

The  first  of  these  transitions  is  the  most  interesting,  since  it  takes  place 
at  ordinary  temperatures.   Below  18°  ordinary  white  or  tetragonal  tin 

500 


TIN  AKD  LEAD  501 

is  an  unstable  form  and  under  some  conditions  changes  into  the  gray 
modification,  at  the  same  time  losing  its  metallic  appearance  and 
crumbling  into  a  powder.  The  usual  form  is  silver-white,  is  harder 
than  lead,  and  is  quite  malleable.  The  bending  of  a  bar  of  tin  is  ac- 
companied by  a  creaking  noise  called  "  tin  cry."  The  metal  melts  at 
232°  and  boils  at  2270°.  Its  average  density  is  7.29. 

Tin  plague.  The  transformation  of  white  tin  into  the  gray  form  is  sometimes 
a  serious  matter.  It  was  first  noticed  in  connection  with  the  deterioration  of 
some  organ  pipes  made  of  tin,  which  developed  holes  and  then  broke  up  com- 
pletely. Later,  ingots  of  tin  stored  in  a  Russian  customhouse  during  a  very 
cold  winter  were  found  to  have  crumbled  to  powder.  When  the  transformation 
once  starts,  it  continues  to  spread  as  long  as  the  temperature  is  below  18°,  and 
this  uncontrollable  spread  has  been  called  the  tin  plague,  or  museum  disease, 
since  it  sometimes  spreads  over  a  tin  object  in  a  museum.  The  speed  of  the 
transformation  is  increased  by  contact  with  a  solution  of  certain  salts  in  which 
tin  is  slightly  soluble. 

Chemical  conduct.  At  ordinary  temperatures  tin  undergoes  prac- 
tically no  change  in  the  air,  nor  is  it  attacked  by  the  combined  action 
of  air  and  water ;  at  higher  temperatures  it  is  covered  with  a  film  of 
oxide.  Dilute  acids  act  upon  it  very  slowly,  evolving  hydrogen,  con- 
centrated hydrochloric  acid  acting  more  rapidly.  Oxidizing  acids, 
such  as  nitric  acid,  convert  it  into  a  hydrate  of  the  oxide  SnO2. 
Under  some  conditions  the  metal  assumes  a  passive  state  resembling 
that  of  iron. 

Uses  of  tin.  Tin  finds  two  great  uses  in  the  industries ;  namely,  the 
manufacture  of  tin  plate  and  of  alloys.  Tin  plate  is  made  by  dipping 
sheets  of  iron  or  steel  into  melted  tin  and  rolling  them  to  uniform 
thickness.  A  similar  coating  is  sometimes  put  upon  sheet  copper  by 
wiping  the  melted  tin  onto  the  clean  surface  with  tow.  A  great  deal 
of  tin  is  now  recovered  annually  from  old  cans.  The  alloys  which 
contain  tin  are  very  numerous.  The  composition  of  some  of  the 
important  ones  is  indicated  in  the  tables  on  pages  375  and  480. 

Compounds  of  tin.  In  its  compounds  tin  is  either  divalent  or  tetra- 
valent,  giving  rise  to  two  series  of  compounds,  known  as  the  stannous 
and  the  stannic.  In  the  stannous  compounds  tin  acts  for  the  most 
part  as  a  base-forming  element,  its  salts  resembling  those  of  zinc  in  a 
general  way. 

As  in  the  case  of  zinc,  the  hydroxide  (Sn(OH)2)  is  soluble  in  strong 
bases,  giving  a  series  of  compounds  known  as  stannites.  As  a  tetra- 
valent  element,  tin  is  chiefly  acid-forming,  its  salts  resembling  those 


502  'GENEBAL  CHEMISTRY  • 

of  silicon.  There  are  also  some  salts  derived  from  tin  as  a  tetravalent 
base,  such  as  the  sulfate  Sn(SO4)2,  as  well  as  a  great  many  double 
and  complex  compounds. 

Stannous  compounds.  Quite  a  number  of  stannous  salts  have  been 
prepared  by  dissolving  either  tin  or  stannous  oxide  in  the  appropriate 
acid.  For  the  most  part  they  are  soluble,  colorless  salts,  usually  form- 
ing a  number  of  hydrates,  and  are  subject  to  considerable  hydrolysis, 
yielding  basic  salts.  Few  of  them  require  special  description. 

Stannous  oxide  (SnO).  This  compound  is  a  black  powder  obtained 
by  warming  stannous  chloride  (SnCl2)  with  sodium  carbonate  and 
thoroughly  washing  the  product  with  hot  water.  The  corresponding 
hydroxide  is  not  known  in  pure  condition,  but  gives  a  dehydration 
product  of  the  composition  2  SnO  •  H2O,  which  in  turn  easily  loses  its 
water  and  is  converted  into  the  oxide.  The  precipitated  hydroxide  is 
soluble  in  alkalies  forming  stannites,  as  shown  in  the  equations 

SnCl2  4-  2  KOH  =  Sn(OH)2  +  2  KC1 
Sn(OH)2  4-  2  KOH  =  K2SnO2  +  2  H2O 

Stannous  sulfide  (SnS).  As  prepared  by  precipitation,  stannous  sul- 
fide  is  a  dark  brown  powder  obtained  by  the  action  of  hydrogen 
sulfide  upon  a  soluble  stannous  salt: 

SnCL  +  HS  =  SnS  +  2  HC1 


•2 


It  is  insoluble  in  water,  in  dilute  acids,  and  in  ordinary  ammonium 
sulfide,  but  is  soluble  in  ammonium  polysulfide  forming  ammonium 
sulf  ostannate : 

SnS  +  (NH4)4S2=(NH4)2SnS3 

Stannous  halides.  The  four  stannous  halides  are  all  prepared  by 
the  usual  methods,  stannous  chloride  (SnCl2)  being  the  best  known. 
The  anhydrous  salt  is  obtained  by  conducting  hydrogen  chloride  over 
heated  tin.  It  is  a  white,  crystalline  compound,  which  melts  at  249° 
and  boils  at  620°.  The  most  familiar  hydrate  has  the  composition 
SnCl2  •  2  H2O  and  is  called  tin  salt.  It  is  used  in  the  dyeing  industry 
as  a  mordant  and  also  as  a  reducing  agent,  since  it  tends  to  pass 
very  readily  to  the  tetravalent  condition.  This  tendency  is  illustrated 
in  the  reaction  with  mercuric  chloride,  as  shown  in  the  equations 

SnCl2  +  2  HgCl2  =  SnCl4  +  2  HgCl 
SnCl2  +  2  FeCl8  =  SnCl4  +  2  FeCl 


TIN  AND  LEAD  503 

Stannic  compounds.  In  its  tetravalent  compounds  tin  usually  plays 
the  part  of  an  acid-forming  element,  the  oxide  SnO2  being  essentially 
an  acid  anhydride,  like  silicon  dioxide  SiO2.  Most  of  these  compounds 
are  colorless  and  are  easily  hydrolyzed. 

Stannic  halides.  With  the  exception  of  the  fluoride,  the  stannic 
halides  are  easily  melted  compounds  of  low  boiling  point,  and  have 
the  general  physical  properties  of  the  halides  of  acid-forming  elements. 
They  are  best  prepared  by  acting  upon  the  metal  with  an  excess  of  the 
halogen.  The  chloride  will  serve  as  a  type  of  these  compounds.  It  is 
a  colorless  liquid  boiling  at  114.1°  and  melting  at  —  33°.  It  dissolves 
in  water  with  strong  heat  evolution  and  with  considerable  contraction 
in  volume.  From  the  solution  five  different  hydrates  have  been  pre- 
pared, the  most  common  of  which  is  the  pentahydrate  SnCl4  •  5  H2O. 
The  chloride  also  combines  directly  with  many  other  compounds,  such 
as  alcohol  and  ammonia.  Its  solution  in  hydrochloric  acid  yields  the 
chloro-acid  H2SnCl6,  and  the  corresponding  fluo-acid  H2SnF6  is  also 
well  known.  These  acids  yield  a  long  series  of  salts,  the  best  known 
of  which  is  ammonium  chlorostannate  ((NH4)2SnCl6),  called  pink  salt, 
which  finds  extensive  use  as  a  mordant  in  the  dyeing  industry. 

Stannic  sulfide  (SnS2).  Stannic  sulfide  is  precipitated  as  a  bright 
yellow,  amorphous  powder,  when  hydrogen  sulfide  is  conducted  into 
an  acid  solution  of  a  stannic  compound.  It  is  insoluble  in  water 
and  in  dilute  acids,  but  is  soluble  in  ammonium  sulfide  forming  a 
sulfostannate:  + 


It  can  be  obtained  in  the  form  of  golden-bronze  scales,  which  feel 
greasy  to  the  touch,  like  graphite,  by  heating  a  mixture  of  stannous 
sulfide,  sulfur,  and  ammonium  chloride.  This  was  formerly  used  as 
a  pigment  called  mosaic  gold. 

Stannic  oxide  (Sn02).  Stannic  oxide  is  the  form  in  which  tin  is 
usually  found  in  nature.  It  occurs  in  three  crystalline  modifications, 
and  when  pure  these  are  colorless  and  transparent.  The  melting  point 
is  quite  high  (1130°).  As  prepared  by  burning  tin  in  air  it  is  an 
amorphous  powder  called  flowers  of  tin.  The  ignited  oxide  is  not 
easily  attacked  by  most  reagents. 

Hydrates  of  stannic  oxide.  The  hydrate  Sn(OH)4,  corresponding 
to  the  dioxide,  is  not  known  in  well-defined  condition.  It  loses  water, 
and  from  it  are  derived  two  distinct  compounds,  which  have  the  same 
percentage  composition,  corresponding  to  the  formula  H2SnO3. 


504  GENERAL  CHEMISTRY 

1.  Stannic  acid.  One  of  these,  called  stannic  acid,  is  obtained  by 
treating  stannic  chloride  with  a  soluble  base :  .1  ^^  Q 

SnCl4  +  4  KOH  =  4  KC1  +  SnO(OH)a  +  H2O 
It  is  soluble  in  excess  of  the  base,  forming  salts  called  stannates  : 
Sn$(pH$g  +  2  KOH  =  K2SnO3  +  2  H2O 

2.  Metastannic  acid.    Metastannic  acid  was  discovered  by  Berzelius 
in  1811,  and  was  the  first  observed  example  of  isomerism.    It  is  pre- 
pared as  a  white,  insoluble  solid  by  oxidizing  tin  with  nitric  acid.    It 
is  less  soluble  than  stannic  acid,  both  in  acids  and  in  alkalies,  and 
gives  salts  of  the  formula  M2Sn6Ou,  such  as  Na2SngOn  •  4  H2O.    This 
makes  it  appear  probable  that  the  free  acid  has  the  formula  (H  SnO  ) 


H10Sn5015. 


LEAD 


History  and  occurrence.  Articles  made  of  lead  have  been  found  in 
Egyptian  ruins  of  great  antiquity,  and  there  is  no  doubt  that  metallic 
lead  found  applications  from  very  early  times.  The  Romans  called  it 
plumbum  and  used  it  for  water  conduits  as  we  do  to-day.  It  does  not 
occur  to  any  appreciable  extent  in  the  native  state,  but  is  found  in 
combination  in  many  parts  of  the  world.  The  minerals  of  most  impor- 
tance to  metallurgists  are  galena,  or  galenite  (PbS),  and  to  a  less 
extent  cerussite  (PbCO8)  and  anglesite  (PbSO4).  The  United  States 
produces  over  one  third  of  the  world's  output  of  lead,  the  chief  pro- 
ducing states  being  Missouri,  Idaho,  Utah,  and  Colorado.  The  coun- 
tries which  come  next  to  the  United  States  in  production  are  Spain, 
Mexico,  and  Germany. 

Metallurgy.  Many  lead  ores  contain  some  silver  and  gold,  and  silver 
ores  are  often  purposely  combined  with  lead  ores  and  the  two  smelted 
together.  The  method  followed  depends  upon  whether  or  not  the  ore 
is  to  be  worked  for  silver.  When  no  silver  is  present,  the  ores  are 
roasted  in  an  open  oven  until  they  are  partially  oxidized,  and  the  lead 
is  present  as  a  mixture  of  sulfide,  sulfate,  and  oxide.  Access  of  air  is 
then  cut  off  and  the  temperature  raised,  when  the  reactions  repre- 
sented in  the  following  equations  take  place  : 


PbS  +  PbSO4  =  2  Pb  +  2  SO2 

Silver-bearing  ores  are  worked  in  a  blast  furnace,  the  bottom  of 
which  consists  of  a  large  crucible  constructed  of  fire  brick.    The  ore 


TIN  AKD  LEAD  505 

is  first  roasted  and  is  then  charged  into  the  furnace,  together  with  coke 
and  a  flux  consisting  of  limestone  and  iron  ore.  The  chief  reactions 
which  produce  lead  are  represented  in  the  equations 

PbS  +  FeO  +  C  -  Pb  +  FeS  +  CO 
PbSO4  +  FeO  +  5  C  =  Pb  +  FeS  +  5  CO 

The  liquid  lead,  in  which  are  dissolved  the  gold  and  silver,  together 
with  varying  quantities  of  copper,  antimony,  arsenic,  and  bismuth, 
collects  in  the  crucible  and  is  tapped  off. 

Refining  of  lead.  The  lead  obtained  in  this  way  is  called  hard  lead 
because  of  the  effect  of  the  alloyed  metals.  It  is  softened  by  the  re- 
moval of  these  as  follows :  The  hard  lead  is  melted  in  a  reverberatory 
furnace  with  free  access  of  air,  -until  the  copper,  arsenic,  and  antimony 
are  oxidized,  together  with  a  considerable  quantity  of  lead.  The  oxides 
are  skimmed  off  and  the  softened  lead,  which  contains  the  silver,  gold, 
and  also  bismuth,  is  run  off  for  desilverizing.  Two  processes  for  this 
purpose  are  in  use  in  this  country. 

1.  The  Parkes  process.  In  the  Parkes  process  the  lead  is  run  into 
kettles  holding  as  much  as  30  tons,  and  about  1  per  cent  of  its  weight 
of  zinc  is  added  and  thoroughly  stirred  in.   These  two  metals  do  not 
mix  to  any  great  extent,  and  gold  and  silver,  as  well  as  copper,  are 
much  more  soluble  in  zinc  than  in  lead.    Consequently,  when  the  stir- 
Ting  ceases,  the  zinc,  together  with  most  of  the  precious  metals,  rises 
to  the  top  and,  when  the  melt  is  allowed  to  cool,  hardens  to  a  crust 
which  can  be  skimmed  off.    The  process  is  repeated  several  times. 
The  zinc  remaining  in  the  lead  is  removed  by  blowing  dry  steam  and 
air  through  the  liquid  lead,  which  oxidizes  the  zinc ;  or  it  is  oxidized 
as  in  the  original  softening.    The  zinc  crusts  are  distilled  from  a  re- 
tort, by  which  process  the  zinc  is  recovered,  to  be  used  again,  and  the 
residue  is  cupeled  as  with  silver  to  give  dore  bars. 

2.  The  Betts  process.  In  the  Betts  process  the  lead  is  refined  by 
electrolysis,  as  in  the  case  of  copper.    A  sheet  of  pure  lead  serves  as 
cathode,  a  thick  plate  of  crude  lead  as  anode,  and  a  solution  of  lead 
fluosilicate  (PbSiF6),  together  with  some  colloidal  material,  such  as 
gelatin,   as  electrolyte.     The   lead  deposits   upon   the    cathode,   the 
iron  remains  dissolved  in  the  electrolyte,  and  the  copper,  bismuth, 
antimony,  arsenic,  silver,  and  gold  are  left  undissolved  as  a  skeleton 
of  the  anode.    A  considerable  quantity  of  bismuth  is  recovered  from 
this  source. 


506  GENERAL  CHEMISTRY 

Properties  of  lead.  Pure  lead  is  a  silvery  metal  of  density  11.37, 
which  melts  at  327°  and  boils  at  1525°.  It  is  dimorphous,  crystalliz- 
ing either  in  the  regular  or  in  the  monoclinic  system.  It  is  the  softest 
of  all  the  heavy  metals  and  is  a  moderately  good  conductor  of  elec- 
tricity. It  is  quite  malleable,  but  has  little  strength. 

Chemical  conduct.  Lead  is  a  moderately  active  metal,  standing  next 
above  hydrogen  .in  the  electromotive  series.  Its  true  activity  is  often 
concealed  by  the  fact  that  so  many  of  its  compounds  are  insoluble 
and  form  protective  coatings  upon  its  surface.  It  quickly  tarnishes 
in  air,  owing  to  the  formation  of  a  bluish-gray  oxide  (Pb2O).  It  is 
acted  upon  with  vigor  by  fluorine,  and  with  less  intensity  by  the 
other  halogens.  It  liberates  hydrogen  from  acids  very  slowly,  and 
apparently  reacts  very  slightly  with  water,  the  reaction  soon  ceasing. 
Oxidizing  acids  attack  it  readily. 

Uses  of  lead.  The  industrial  uses  of  lead  are  very  numerous.  Chief 
among  them  are  the  manufacture  of  water  pipes,  of  storage  batteries, 
and  of  structures  to  be  exposed  to  acids,  such  as  the  lead  chambers  in 
a  sulfuric  acid  plant.  A  great  many  alloys  contain  lead,  among  them 
being  type  metal,  antifriction  metals  of  various  kinds,  pewter,  and 
solder  (p.  375).  Over  one  third  of  the  yearly  output  is  used  in  the 
manufacture  of  paints  and  is  permanently  lost. 

Compounds  of  lead.  With  very  few  exceptions  lead  is  either  divalent 
or  tetravalent  in  its  compounds.  The  divalent  hydroxide  (Pb(OH)2) 
is  essentially  a  base  and  gives  rise  to  a  series  of  well-defined  salts, 
most  of  which  are  colorless,  since  lead  gives  them  no  characteristic 
color.  They  are  somewhat  hydrolyzed  in  solution  and  yield  numerous 
basic  salts.  The  tetravalent  hydroxide  (Pb(OH)4)  is  an  acid,  and  but 
few  of  its  derivatives  are  well  denned.  All  compounds  of  lead  which 
are  at  all  soluble  are  poisonous,  and  lead  workers  of  all  classes  are 
subject  to  the  occupational  disease  known  as  lead  colic. 

Oxides  of  lead.  Lead  forms  five  oxides,  of  which  three  are  simple 
oxides  having  the  formulas  Pb2O,  PbO,  and  PbO2.  The  other  two, 
whose  formulas  are  Pb2O8  and  Pb3O4,  are  complex  compounds. 

Lead  suboxide  (Pb20).  This  compound  is  obtained  as  a  grayish-black 
powder  by  carefully  heating  lead  oxalate : 

2  PbC204  =  Pb20  +  3  C02  +  CO 

It  forms  as  a  thin  film  on  the  surface  of  exposed  lead,  to  which  it  gives 
the  characteristic  lead  color. 


TIN  AND  LEAD  507 

Lead  monoxide  (litharge)  (PbO).  This  oxide  appears  to  exist  in  a 
number  of  modifications,  the  colors  of  which  range  from  yellow  and 
light  brown  to  red,  and  the  color  of  the  commercial  product  varies 
considerably  in  consequence.  It  is  a  highly  crystalline  compound,  but 
commercial  litharge,  obtained  as  a  by-product  in  a  number  of  processes, 
is  usually  a  fine  powder.  The  corresponding  hydroxide  (Pb(OH)2) 
is  a  white  solid  obtained  by  precipitation.  It  is  very  slightly  soluble 
in  water,  and  in  the  process  of  drying  forms  two  dehydration  products. 
It  is  soluble  in  acids,  giving  the  best-known  salts  of  lead,  and  in  the 
strong  alkalies,  forming  compounds  called  plumbites : 

Pb(OH)2  +  2  KOH  =  K2PbO2  +  2  H2O 

Lead  dioxide  (Pb02).  This  compound,  very  frequently  called  lead 
peroxide,  is  prepared  by  the  action  of  chlorine  or  sodium  hypochlo- 
rite  upon  an  alkaline  solution  of  lead  acetate.  It  is  also  obtained  as  a 
coating  upon  the  anode  when  solutions  of  lead  salts  are  subjected  to 
electrolysis.  It  is  a  chocolate-brown  powder  and  is  a  good  oxidizing 
agent.  Thus,  with  hydrochloric  acid  it  acts  in  a  manner  similar  to 
manganese  dioxide,  liberating  chlorine : 

Pb02  +  4  HC1  =  PbCl2  +  C12  +  2  H20 

It  plays  an  important  part  in  the  chemistry  of  storage  cells  (p.  511). 
Dioxides  and  peroxides.  It  will  be  noted  that  the  formulas  of  lead 
dioxide  (PbO2)  and  barium  peroxide  (BaO2)  are  very  similar.  The  lat- 
ter of  these  yields  hydrogen  dioxide  when  treated  with  an  acid,  and  it 
is  regarded  as  a  compound  in  which  an  atom  of  barium  has  displaced 
two  atoms  of  hydrogen  in  hydrogen  peroxide,  and  consequently  as 

having  the  structure  Ba<^  I  .    All  the  salts  of  hydrogen  peroxide  are 
known  as  peroxides. 

Lead  dioxide  and  the  similar  manganese  dioxide  (MnO2)  yield  no 
hydrogen  peroxide  when  treated  with  acids,  but  act  as  oxidizing 
agents,  being  changed  into  derivatives  of  the  lower  oxides  (PbO  and 
MnO).  The  metal  in  these  compounds  is  regarded  as  tetravalent,  so 
that  their  structure  is  represented  by  the  formulas  Pb^o  and  Mn^^. 
The  term  dioxide  is  therefore  applied  to  oxides  of  tetravalent  elements, 
in  order  to  distinguish  them  from  salts  of  hydrogen  dioxide. 

Lead  nitrate  (Pb(N03)2).  This  salt  is  readily  obtained  by  dissolving 
metallic  lead  or  litharge  in  nitric  acid.  It  crystallizes  in  octahedra 
and  is  easily  soluble  in  water. 


508  GENERAL   CHEMISTRY 

Lead  acetate  (Pb(C2H302)2).  The  acetate  is  obtained  by  dissolving 
litharge  in  concentrated  acetic  acid,  from  which  it  crystallizes  in  snow- 
white,  monoclinic  crystals  of  the  composition  Pb(C2HgO0)2  •  3  H2O. 
It  is  known  as  sugar  of  lead  because  of  its  sweetish  taste.  Several 
soluble  basic  acetates  are  known,  the  most  familiar  one  having  the 
composition  expressed  by  the  formula  Pb(C2H3O2)OH.  The  acetate 
and  nitrate  are  the  important  soluble  salts  of  lead. 

Lead  sulfide  (PbS).  Native  lead  sulfide  (galena)  has  almost  the 
appearance  of  lead  itself,  save  that  it  is  conspicuously  crystalline. 
Prepared  by  precipitation,  it  is  a  black,  amorphous  solid,  insoluble 
in  water  and  in  acids. 

The  halides  of  lead.  Lead  chloride  (PbCl2)  is  precipitated  as  a 
white,  crystalline  powder  when  a  soluble  lead  salt  is  treated  with  a 
soluble  chloride.  It  is  very  sparingly  soluble  in  cold  water  and  in 
acids,  but  is  more  soluble  in  hot  water.  The  bromide  (PbBr2)  and 
the  iodide  (PbI2)  resemble  the  chloride  in  properties,  save  that  the 
iodide  is  golden  yellow  in  color.  A  tetrachloride  (PbCl4)  correspond- 
ing to  the  dioxide  has  been  prepared  in  the  following  way :  The  di- 
chloride  is  dissolved  in  hydrochloric  acid,  and  chlorine  is  conducted 
into  the  very  cold  solution.  Ammonium  chloride  is  then  added,  and 
a  crystalline  compound  ((NH4)2PbCl6)  is  obtained.  This  is  treated 
with  cold  concentrated  sulf  uric  acid,  when  a  reaction  takes  place  which 
liberates  the  tetrachloride  as  a  heavy,  yellowish  liquid  resembling  tin 
tetrachloride : 

(NH4)2PbCl6  +  H2S04  =  (NH4)2S04  +  2  HC1  +  PbCl4 

Lead  carbonate  (PbC03).  Normal  lead  carbonate,  which  occurs  in 
nature  as  cerussite,  may  be  prepared  as  a  white,  crystalline  powder 
by  precipitating  a  solution  of  a  lead  salt  with  a  solution  of  sodium 
carbonate.  Several  basic  carbonates  of  lead  are  known,  but  the  one 
having  the  formula  2  Pb(CO8)2  •  Pb(OH)2,  called  white  lead,  is  of 
much  technical  importance,  since  it  is  the  basis  of  most  paints. 

Manufacture  of  white  lead.  White  lead  can  be  prepared  by  a  number  of 
processes,  but  none  of  them  seems  to  produce  a  product  of  as  desirable  physical 
properties  as  the  old  Dutch  process,  which  has  been  used  for  centuries,  though 
with  many  improvements.  In  this  process  the  lead  is  cast  into  perforated  plates, 
which  are  placed  loosely  upon  each  other  in  a  crock  of  the  shape  shown  in 
Fig.  152,  the  ledge  formed  by  the  constriction  in  the  crock  supporting  the  plates. 
Under  them  is  poured  a  suitable  quantity  of  dilute  acetic  acid,  and  the  crocks 
so  charged  are  placed  in  banks  and  covered  with  stable  manure  or  spent  tanbark. 


TEST  AND  LEAD 


509 


FIG.  152 


The  heat  of  fermentation  in  the  latter  warms  the  acid,  the  fumes  of  which  attack 
the  lead,  forming  acetate.  The  carbon  dioxide  from  the  fermentation  enters  into 
reaction  with  the  acetate  and  produces  the  basic  carbonate,  regenerating  acetic 
acid,  which  acts  again  upon  the  lead.  The  process  con- 
tinues until  the  plates  are  almost  completely  converted 
into  the  desired  compound. 

Paints.  The  manufacture  of  paints  is  a  very  exten- 
sive chemical  industry,  and  absorbs  a  large  percentage 
of  all  the  lead  produced.  A  paint  consists  of  three 
essential  ingredients  : 

1.  The  vehicle,  or  liquid  medium.    This  must  be  an  oil 
which  will  dry  rapidly  and  harden  in  drying  to  a  more 
or  less  flexible,  hornlike  body.    These  changes  in  the 
oil  are  due  to  oxidation  by  the  air.    A  number  of  dif- 
ferent oils  will  serve  this  purpose,  but  linseed  oil  has 
long  been  used  as  the  standard  drying  oil,  since  it  can 
be  produced  in  quantity  and  at  moderate  cost.    It  is 

customary  to  add  to  it  a  dryer,  made  by  boiling  some  of  the  oil  with  oxides  of 
manganese,  lead,  or  cobalt.  The  oxides  enter  into  combination  with  the  oil  and 
assist  catalytically  in  its  oxidation. 

2.  The  body.  The  body  of  the  paint  must  be  some  solid  material,  suspended  in 
the  oil,  which  will  give  a  smooth  and  waxy  surface  as  the  paint  dries,  and  will 
have  good  covering  power.    While  white  lead  meets  these  requirements,  it  is 
moderately  expensive  and  also  blackens  when  exposed  to  hydrogen  sulfide,  which 
is  likely  to  be  present  in  the  air  in  cities.    Other  bodies  are  now  frequently  com- 
bined with  the  lead  or  replace  it  altogether,  among  them  being  zinc  oxide,  China 
clay  (or  kaolin),  barium  sulfate,  and  a  product  called  lithophone.    This  is  a  com- 
bination of  zinc  sulfide  and  barium  sulfate  produced  by  precipitating  barium 
sulfide  with  zinc  sulfate  : 

BaS  +  ZnSO4  =  BaSO4  +  ZuS 

For  some  purposes  these  materials  are  a  real  advantage,  and  they  are  not  to  be 
regarded  as  adulterants  unless  sold  as  white  lead. 

3.  The  pigment,  or  coloring  matter.  In  the  case  of  white  paints  the  body  serves 
also  as  the  coloring  matter.    For  other  colors  a  specific  pigment  must  be  added. 
In  most  cases  these  are  metallic  oxides  or  salts,  and  are  frequently  natural  products. 
Sometimes  they  are  prepared  by  precipitating  an  amorphous  body  (usually  a 
colloid)  in  the  presence  of  an  organic  dye,  the  dye  being  absorbed  by  the  precipi- 
tate and  giving  it  a  color.    Such  pigments  can  be  prepared  in  endless  variety  of 
color  and  are  called  lakes.    They  are  not  usually  as  permanent  as   a  mineral 
pigment. 

Lead  sulfate  (PbSOJ.  Lead  sulfate  is  a  white,  crystalline  solid, 
insoluble  in  water  and  dilute  acids.  It  is  therefore  formed  as  a  pre- 
cipitate whenever  the  ions  Pb+  +  and  SO4~ "  are  brought  together 
in  solution.  The  sulfate  is  soluble  in  concentrated  sulfuric  acid,  but 
precipitates  again  when  this  is  diluted. 


510  GENERAL   CHEMISTRY 

Lead  chromate  (PbCrOJ.  This  bright  yellow  salt,  called  chrome 
yellow,  results  as  a  precipitate  when  a  soluble  lead  salt  is  treated  with 
a  soluble  chromate : 

Pb(CfH.O,),  +  K,Cr04  =  PbCrO,  +  2  KC.H.O, 

By  boiling  the  normal  chromate  with  a  solution  of  an  alkali,  a  brick- 
red  basic  salt  is  obtained,  called  chrome  red,  which  has  the  formula 
PbCrO4  •  PbO.  Both  of  these  chromates  are  used  as  paint  pigments. 

Lead  arsenate  (Pb3(As04)2).  This  compound  is  a  white,  insoluble 
powder  difficultly  soluble  in  water,  prepared  by  treating  lead  acetate 
with  sodium  arsenate.  It  is  quite  extensively  employed  as  an 
insecticide. 

Acids  of  lead.  Both  of  the  hydroxides  of  lead  have  feeble  acid 
properties,  dissolving  in  solutions  of  the  alkalies  to  form  salts. 

The  plumbites.  The  salts  derived  from  plumbous  hydroxide  acting 
as  an  acid  are  called  plumbites.  The  formation  of  sodium  plumbite  is 
represented  in  the  equation 

Pb(OH)2  +  2  NaOH  =  Na2PbO2  +  2  H2O 

The  plumbates.  The  derivatives  of  plumbic  hydroxide,  acting  as  an 
acid,  are  called  plumbates.  Some  of  these  are  orthoplumbates,  derived 
from  orthoplumbic  acid  (H4PbO4),  while  others  are  metaplumbates, 
derived  from  metaplumbic  acid  (H2PbO3). 

Lead  orthoplumbate  (minium,  or  red  lead)  (Pb3OJ.  The  formula 
PbgO4  is  usually  assigned  to  this  salt,  though  its  composition  is  more 
satisfactorily  expressed  by  the  formula  Pb.2PbO4.  It  is  a  bright  red 
powder  obtained  by  heating  litharge  in  the  air  to  about  450°,  and  is 
valuable  as  a  paint  pigment.  When  treated  with  nitric  acid,  two 
thirds  of  the  lead  passes  into  solution,  the  other  third  remaining  as 
the  insoluble  dioxide : 

Pb304  +  4  HN03  =  2  Pb(N03)2  +  PbO2  +  2  H2O 

Calcium  orthoplumbate  (Ca2PbO4)  is  obtained  by  heating  a  mixture 
of  lead  dioxide  and  calcium  oxide. 

Lead  metaplumbate  (Pb203).  Among  the  metaplumbates,  one  of  the 
most  interesting  is  the  compound  Pb2O3.  It  is  an  orange-yellow 
powder,  the  composition  of  which  is  more  satisfactorily  expressed  by 
the  formula  PbPbO3.  Nitric  acid  dissolves  one  half  of  the  lead  from 
this  compound,  leaving  the  other  half  as  lead  dioxide. 


TIN  AND  LEAD 


511 


FIG.  153 


Storage  cell.  The  storage  cell,  or  accumulator,  plays  an  important  part  in 
modern  electrical  developments.  Its  fundamental  characteristic  is  that  the 
chemical  action  upon  which  it  depends  is  revers- 
ible. The  chemical  action  taking  place  when 
the  cell  is  delivering  current  is  reversed  when  a 
current  is  conducted  through  the  cell  in  an 
opposite  direction.  Electrical  energy  can  there- 
fore be  stored  in  the  cell  as  chemical  energy  and 
drawn  off  again,  when  desired,  as  electrical 
energy.  Many  chemical  reactions  are  adapted 
to  this  purpose,  including  those  which  take  place 
in  the  common  Daniell  cell  (p.  485),  but  there 
are  a  great  many  purely  physical  and  mechanical 
requirements  which  are  difficult  to  meet,  and  in 
practice  only  two  types  of  cells  have  proved  suc- 
cessful. These  are  usually  known  as  the  chloride 
accumulator  and  the  Edison  cell.  . 

In  the  accumulator  (Fig.  153)  the  electrodes 
are  made  of  a  skeleton  of  lead.  When  ready  for 
use,  the  one  plate  is  covered  with  a  thick  de- 
posit of  spongy  lead,  which  is  the  active  mate- 
rial ;  the  other  is  similarly  covered  with  a  layer 

of  lead  dioxide.  The  electrolyte  is  moderately  dilute  sulfuric  acid.  A  number 
of  pairs  of  such  plates  are  arranged  together  in  one  cell.  When  the  plates  are 
connected  by  a  wire,  the  reactions  are  as  follows  : 

At  the  lead  plate  : 

Pb  +  SO4~-  =  PbSO4  +  2  (-) 

The  insoluble  lead  sulfate  deposits  in  the  spongy  lead,  and  the  negative  charge 
is  given  up  to  the  plate. 
At  the  lead  dioxide  plate  : 

Pb02  +  H2S04  +  2  H+  =  PbS04  +  2  H20  +  2  (  +  ) 

The  lead  sulfate  deposits  with  the  dioxide,  and  the  positive  charge  is  given  up 
to  the  plate.  The  complete  equation  is  therefore 

Pb  +  PbO2  +  2  H2SO4  =  2  PbSO4  +  2  H2O 

It  will  be  seen  that  the  action  of  the  cell  results  in  bringing  the  two  plates  to  an 
identical  condition  and  in  withdrawing  sulfuric  acid  from  the  electrolyte.  The 
cell  is  never  allowed  to  come  entirely  to  this  discharged  condition.  When  the 
current  is  reversed,  the  two  plates  are  restored  to  their  original  state. 

In  the  Edison  cell  the  one  plate  is  of  iron,  the  other  is  covered  with  a  deposit 
of  the  higher  oxide  of  nickel  (Xi2O3),  and  the  electrolyte  is  a  solution  of  potas- 
sium hydroxide.  The  reactions  are  not  so  well  understood  as  in  the  case  of 
the  lead  battery,  but  in  a  general  way  the  following  equation  represents  the 
reversible  reaction : 


Fe 


2  Xi(OH)2  +  Fe(OH)2 


512  GENERAL  CHEMISTRY 

GERMANIUM 

Germanium  was  discovered  in  1886  by  Winkler,  in  connection  with 
some  analyses  of  the  rare  mineral  argyrodite.  His  analyses  failed  to 
account  for  from  6  to  7  per  cent  of  the  mineral,  and  a  long  and  care- 
ful search  resulted  in  the  discovery  of  the  new  element.  It  was  found 
•to  fulfill  in  a  remarkable  way  the  predictions  of  Mendeleeff  for  the 
properties  of  an  undiscovered  element  which  should  follow  silicon  in 
the  fourth  group,  and  which  he  provisionally  named  ekasijicon. 

In  the  elementary  condition,  germanium  is  a  soft,  crystalline,  metallic 
substance  having  a  density  of  5.47  and  melting  at  about  900°.  It 
forms  two  oxides  of  the  formulas  GeO  and  GeO2,  each  of  which  gives 
rise  to  a  series  of  compounds.  Those  in  which  germanium  is  tetrava- 
lent  recall  the  corresponding  compounds  of  carbon  and  silicon.  Those 
derived  from  the  lower  oxide  ha"ve  been  very  little  investigated. 

THORIUM 

Thorium,  named  in  honor  of  the  Scandinavian  god  Thor,  was  dis- 
covered by  Berzelius  in  1828.  The  element  is  an  essential  constituent 
of  a  few  rare  minerals,  notably  thorite  (ThSiO4),  but  for  the  most 
part  it  is  found  in  very  small  concentrations  in  various  minerals,  espe- 
cially in  those  which  are  rich  in  the  rare  earths  (p.  451).  Commer- 
cially its  compounds  are  almost  entirely  obtained  from  monazite  sand, 
which  is  essentially  a  phosphate  of  the  rare  earths,  but  which  carries 
from  0  to  8  per  cent  thorium  phosphate. 

The  metal  is  very  difficult  to  prepare  in  pure  condition,  and  is  best 
obtained  by  reducing  the  oxide  with  mixed  metal  (p.  452).  It  is  a 
heavy  metal  somewhat  resembling  platinum  in  luster,  hardness,  and 
ductility.  It  melts  at  1690°.  In  its  compounds  thorium  always  acts 
as  a  tetravalent  metal,  and  it  forms  a  long  series  of  simple  salts,  as 
well  as  many  double  ones.  The  chief  salt  of  commerce  is  the  nitrate, 
the  usual  hydrate  of  which  has  the  formula  Th(NO3)4  •  12  H2O.  This 
salt  finds  an  extensive  use  in  the  manufacture  of  gas  mantles  of  the 
Welsbach  type. 

Gas  -mantles.  A  gas  mantle  is  essentially  an  envelope,  of  very  small  weight  and 
large  surface,  which  is  suspended  over  a  nonluminous  Bunsen  flame  and  which 
becomes  brilliantly  luminous  at  a  relatively  low  temperature.  The  physical 
requirements  for  such  a  material  «ire  very  exacting.  For  a  given  weight  it  must 
be  workable  into  a  large,  porous  surface ;  it  must  become  luminous  at  a  low  tem- 
perature ;  it  must  be  efficient  for  a  considerable  period ;  it  must  have  a  very 


TENT  AND  LEAD  513 

small  coefficient  of  expansion,  so  as  to  withstand  sudden  and  severe  changes  in 
temperature,  and  it  must  not  be  too  fragile.  A  mixture  of  99  per  cent  thorium 
oxide  (ThO2)  and  1  per  cent  cerium  oxide  (CeO2)  has  been  found  to  meet  all  of 
these  requirements  admirably.  The  mantle  is  prepared  by  dipping  a  cotton  wick, 
woven  in  the  desired  shape,  into  a  solution  containing  the  nitrates  of  thorium 
and  cerium  in  the  proper  proportion.  The  wick  is  then  dried  on  a  form  and 
very  carefully  burned.  In  this  process  the  nitrates  are  converted  into  oxides, 
which  retain  the  form  of  a  porous  skeleton  of  the  original  cotton  wick.  This  is 
dipped  into  collodion,  to  render  it  less  fragile  in  handling,  the  collodion  being 
burned  off  again  when  the  mantle  is  in  position  on  the  gas  jet.  To  supply  the 
thorium  for  this  industry,  many  tons  of  monazite,  chiefly  from  Brazil  and  South 
Carolina,  are  worked  over  each  year. 

Radioactivity  of  thorium.  The  atomic  weight  of  thorium  is  232.4, 
and  this  is  the  largest  possessed  by  any  of  the  elements,  excepting 
uranium  (238.5).  It  is  therefore  extremely  interesting  to  find  that 
compounds  of  thorium,  like  those  of  uranium,  possess  a  peculiar 
property  known  as  radioactivity.  As  this  property  was  discovered  in 
connection  with  the  latter  element,  in  which  it  is  much  more  pro- 
nounced, a  discussion  of  it  will  be  delayed  until  the  compounds  of 
uranium  are  described  (p.  535). 


CHAPTER  XXXYII 

MANGANESE  AND  CHROMIUM 

General.  The  elements  manganese  and  chromium  occur  in  different 
periodic  families,  but  there  are  certain  advantages  in  considering  them 
together.  Neither  one  is  very  closely  related  to  any  other  well-known 
element,  manganese  having  no  companion  element  in  the  seventh 
group  and  chromium  differing  considerably  from  the  other  members 
of  its  family.  On  the  other  hand,  the  two  elements  have  a  good  many 
characteristics  in  common.  In  their  chemical  conduct  manganese  and 
chromium  present  about  the  greatest  variety  of  all  the  elements,  with 
the  exception  of  carbon.  Each  exists  in  a  number  of  stages  of  oxi- 
dation, each  stage  being  represented  by  a  series  of  compounds.  In 
lower  valences  these  elements  are  base-forming,  while  in  higher  valences 
they  are  acid-forming,  and  since  they  pass  readily  from  one  condition 
to  the  other,  with  corresponding  oxidation  or  reduction,  there  is  great 
variety  in  the  reactions  accompanying  these  changes.  In  describing 
the  compounds  of  these  elements  it  will  be  possible  to  mention  only 
those  which  will  serve  to  illustrate  the  characteristics  of  each  series. 

MANGANESE 

History  and  occurrence.  While  manganese  is  not  at  all  a  rare  ele- 
ment, its  ores  were  confused  with  those  of  iron  by  the  earlier  chemists, 
and  it  was  not  until  the  time  of  Scheele  (1774)  that  the  mineral 
pyrolusite  was  shown  to  be  essentially  different  from  magnetite  and 
to  contain  a  different  metal.  Manganese  occurs  in  nature  chiefly  as 
the  dioxide  MnO2,  known  as  pyrolusite.  The  largest  deposits  are  in 
Russia,  India,  and  Brazil ;  in  the  United  States,  Virginia  is  the  chief 
producing  state.  The  element  also  occurs  in  the  form  of  a  number 
of  other  oxides  and  their  hydrates,  and  is  very  widely  distributed  in 
small  percentages  through  many  minerals  and  soils.  To  some  extent 
it  is  absorbed  by  plants. 

Preparation.  Entirely  pure  manganese  is  difficult  to  prepare,  since 
the  metal  tends  to  combine  with  many  of  the  usual  reducing  agents, 
such  as  carbon.  It  is  most  easily  prepared  in  a  fairly  pure  state  by  the 

514 


MANGANESE  AND  CHROMIUM  515 

reduction  of  its  oxide  by  aluminium  (Goldschmidt  method)  or  by  care- 
fully controlled  reduction  by  carbon  in  an  electric  furnace.  It  is  much 
more  frequently  produced  as  an  alloy  by  the  simultaneous  reduction 
of  its  oxide  with  that  of  some  other  metal,  such  as  iron  or  copper. 

Properties  and  conduct.  Manganese  is  a  hard  and  brittle  metal  some- 
what resembling  iron  in  appearance,  but  often  with  a  slightly  reddish 
tint.  It  melts  at  about  1207°  and  boils  at  about  1900°,  both  these 
temperatures  being  considerably  lower  than  the  corresponding  ones 
for  iron.  Its  density  is  7.39. 

In  chemical  conduct  manganese  most  closely  resembles  iron.  It 
oxidizes  in  the  air  with  great  ease  when  pure,  but  less  rapidly  when 
it  contains  some  carbon.  It  liberates  hydrogen  from  dilute  acids  and 
from  water. 

Uses  of  manganese.  The  largest  use  of  manganese  is  as  an  alloy 
constituent  in  steel.  Sometimes  the  manganese  ore  is  added  to  the 
iron  ore  in  the  blast  furnace,  and  many  iron  ores  already  contain  some 
manganese.  In  other  cases  a  rich  alloy  of  manganese  is  first  prepared, 
and  this  is  added  as  may  be  desired.  Ferromanganese  contains  about 
70  per  cent  of  manganese,  while  Spiegel  iron  contains  from  5  to  15  per 
cent.  Manganese  bronze  is  an  alloy  with  copper,  while  manganin  con- 
tains 84  per  cent  of  copper,  4  of  nickel,  and  12  of  manganese.  The 
latter  is  used  as  standard  resistance  wire  in  electrical  measurements. 

Compounds  of  manganese.  Manganese  yields  compounds  corre- 
sponding to  five  different  valences.  The  oxides  from  which  these 
compounds  are  derived  have  the  formulas  MnO,  Mn2O3,  MnO,, 
Mn2O5,  and  Mn0O7,  all  of  which  are  known,  together  with  several  other 
compound  oxides,  such  as  Mn3O4.  When  these  oxides  are  treated  with 
acids,  they  tend  to  give  up  oxygen  and  yield  salts  derived  from  the 
lowest  oxide,  MnO.  If  the  acid  is  one  which  can  be  oxidized,  as  hydro- 
chloric acid,  the  oxygen  is  not  given  up  as  such,  but  goes  to  oxidize 
the  acid.  On  the  other  hand,  when  treated  with  strong  bases  in  the 
presence  of  air,  the  oxides  of  lower  valence  take  up  oxygen  and  yield 
derivatives  of  the  two  highest  oxides.  These  two  tendencies  are  con- 
spicuous not  only  in  the  oxides  themselves  but  in  all  the  compounds 
derived  from  them,  and  most  of  the  transformations  of  the  compounds 
of  manganese  can  be  traced  to  this  broad  principle. 

Manganous  compounds.  The  manganous  compounds  are  derivatives 
of  manganous  oxide  (MnO).  The  oxide  itself  is  a  greenish  powder 
obtained  by  the  reduction  of  the  higher  oxides  with  hydrogen  or  by 


516  GENERAL  CHEMISTRY 

heating  the  carbonate  out  of  contact  with  air.  The  corresponding 
hydroxide  (Mn(OH)2)  is  obtained  as  a  white  precipitate  by  treating  a 
solution  of  any  manganous  salt  with  a  soluble  base.  Under  these  con- 
ditions, however,  it  is  in  contact  with  a  base,  and  it  oxidizes  in  the  air 
with  great  rapidity,  changing  into  hydrates  of  higher  oxides.  The 
salts  derived  from  manganous  oxide  are  quite  stable,  well-crystallized 
compounds.  The  soluble  ones  are  light  pink  in  color,  while  those 
prepared  by  precipitation,  such  as  the  sulfide  (MnS)  and  the  carbonate 
(MnCOg)  are  nearly  white.  All  these  salts  tend  to  form  many  hydrates, 
the  one  obtained  in  any  given  case  depending  upon  the  temperature 
at  which  the  salt  crystallizes.  The  most  familiar  hydrate  of  the  chlo- 
ride has  the  composition  MnCl2  •  4  H2O,  and  the  sulfate  obtained  by 
crystallization  at  room  temperature  is  MnSO4  •  4  H2O,  though  at  some- 
what lower  temperatures  the  hydrate  with  five  as  well  as  with  seven 
molecules  of  water  can  be  obtained. 

Manganic  compounds.  While  a  number  of  manganic  compounds 
have  been  described,  few  of  them  are  well  defined.  The  oxide  (Mn2O3) 
is  obtained  as  a  black  powder  when  manganous  hydroxide  is  exposed 
to  the  air  and  the  product  is  carefully  dried.  It  is  probable,  however, 
that  the  oxide  does  not  contain  trivalent  manganese,  as  its  formula 
would  suggest,  but  that  it  is  a  complex  body,  made  up  as  represented 
by  the  formula  MnO  •  MnO2  and  similar  to  the  corresponding  oxide 
of  lead.  There  is  evidence  of  the  existence  of  a  chloride  (MnCl3) 
and  a  sulfate  (Mn2(SO4)3),  as  well  as  of  a  few  other  trivalent  salts. 
Most  of  these  salts  have  a  deep  cherry-red  color.  None  of  them  play 
an  important  part  in  the  chemistry  of  manganese. 

Manganese  dioxide  (Mn02);  the  manganites.  Manganese  dioxide 
occurs  in  nature  not  only  in  the  form  of  the  hard,  dense  mineral 
pyrolusite  (MnO2)  but  in  a  number  of  hydrated  forms  and  in  combina- 
tion with  other  oxides,  as  in  hausmannite  (MnO2  •  2  MnO).  Owing  to 
the  fact  that  pyrolusite  readily  gives  up  oxygen,  which  is  not  a  very 
common  property  in  minerals,  it  has  long  been  used  as  an  oxidizing 
agent  in  the  industries,  especially  in  the  production  of  chlorine.  It  is  also 
used  in  glass  making  (p.  455).  As  prepared  in  the  laboratory,  it  is  a 
dark  brown  to  black  powder.  It  is  produced  upon  the  anode  when 
manganese  salts  are  subjected  to  electrolysis,  like  the  corresponding 
dioxide  of  lead.  Like  the  latter  compound,  it  acts  as  the  anhydride  of 
a  weak  acid,  and  salts  corresponding  to  the  plumbates  and  stannates 
may  be  obtained  by  heating  the  dioxide  with  the  oxides  of  basic 


MANGANESE  AND  CHROMIUM  517 

elements,  such  as  calcium.  These  salts  are  called  manyanites,  the  formula 
of  the  calcium  salt  being  CaMnO3,  or  CaO  •  MnO2.  The  corresponding 
manganous  salt  is  MnO  •  MnO2,  or  Mn2O3,  which  occurs  in  nature  as 
braunite. 

When  treated  with  very  cold,  concentrated  hydrochloric  acid,  the 
oxide  appears  to  form  a  tetrachloride  (MnCl4)  analogous  to  lead 
tetrachloride.  Like  the  latter  compound,  it  is  very  unstable,  decom- 
posing first  into  the  trichloride  (MnCl3)  and  then  into  the  stable 
dichloride  (MnCl2). 

Manganic  acid  (H2Mn04)  ;  the  manganates.  When  manganese  diox- 
ide or  any  of  the  lower  oxides  is  fused  with  an  alkali  in  the  presence 
of  air  or  an  oxidizing  agent,  the  manganese  passes  to  'the  hexavalent 
state,  and  salts  of  manganic  acid  (H2MnO4)  are  formed  : 

2  MnO2  +  4  KOH  +  O2  =  2  K2MnO4  +  2  H2O 

In  formulas,  and  often  in  crystalline  form,  the  manganates  are  analo- 
gous to  the  sulfates  and  chromates.  The  soluble  sodium  and  potassium 
salts  give  deep  green  solutions  and  very  dark-colored  green  crystals. 
Manganic  acid  is  not  known  in  the  free  state,  but  it  is  evidently  a 
very  much  weaker  acid  than  either  sulfuric  or  chromic  acid.  This  is 
shown  by  the  fact  that  even  its  sodium  and  potassium  salts  are  very 
greatly  hydrolyzed  in  solution  and  are  stable  only  in  the  presence  of 
a  considerable  excess  of  free  base.  When  this  excess  is  constantly 
removed  by  neutralization  with  some  other  acid,  such  as  nitric  or 
even  carbonic  acid,  manganic  acid  is  set  free,  but  it  at  once  undergoes 
a  very  interesting  transformation,  which  is  without  any  analogies 
among  other  well-known  compounds  of  similar  character,  forming  a 
compound  known  as  permanganic  acid. 

Permanganic  acid  (HMnOJ  ;  the  permanganates.  The  formation  of 
permanganic  acid  from  manganic  acid  can  best  be  understood  by 
keeping  in  mind  the  fact  that  the  hexavalent  condition  of  manganese  is 
much  less  stable  than  either  the  heptavalent  or  the  tetravalent  state. 
Accordingly,  when  manganic  acid  is  set  free  it  undergoes  a  rearrange- 
ment whereby  some  of  it  advances  in  valence  to  the  heptavalent  state, 
and  at  the  same  time  some  is  reduced  to  the  tetravalent  form,  the 
transformation  being  expressed  in  the  equation 

3  H2MnO4  =  2  HMnO4  +  Mn(OH)4, 


or  o*       \OH~     0^O         HO/       M>H 


518  GENERAL  CHEMISTRY 

Since  the  permanganic  acid  formed  in  the  reaction  is  a  strong,  soluble 
acid,  it  is  the  salt  and  not  the  free  acid  which  is  obtained.  The  com- 
plete equation  then  becomes 

3  K2Mn04  +  4  H20  =  2  KMnO4  +  Mn(OH)4  +  4  KOH 
The  anhydride  of  permanganic  acid  (Mn2O7)  is  obtained  by  very 
cautiously  adding  concentrated  sulfuric  acid  to  crystals  of  potassium 
permanganate  : 

2  KMn04  +  H2S04  =  K2SO4  +  Mn2O7  +  H2O 

It  is  a  greenish  oil  and  is  an  extremely  energetic  oxidizing  agent.  It 
spontaneously  decomposes,  sometimes  explosively,  forming  the  dioxide 
and  free  oxygen  i  g  =  4  +  g 


Potassium  permanganate  (KMnOJ.  From  the  foregoing  discussion 
it  will  be  seen  that  potassium  permanganate  is  prepared  by  fusing 
potassium  hydroxide,  manganese  dioxide,  and  an  oxidizing  agent 
(usually  potassium  chlorate),  dissolving  the  manganate  so  formed  in 
water,  and  neutralizing  the  excess  of  alkali  with  some  nonoxidizable 
acid.  The  salt  is  very  soluble,  forming  a  deep  reddish-purple  colored 
solution  from  which  it  crystallizes  in  rhombic  needles  of  a  purple- 
black  color  with  a  greenish  reflection.  Other  permanganates  have  the 
same  color  in  solution  as  does  the  free  acid  (HMnO4),  and  those 
derived  from  moderately  strong  bases  are  all  soluble  and  are  little 
hydrolyzed  in  solution.  The  free  acid  can  be  prepared  by  electrolysis 
of  the  potassium  salt,  and  from  this  the  other  salts  can  be  made. 

Oxidizing  action  of  potassium  permanganate.  Potassium  perman- 
ganate finds  extensive  use  as  an  oxidizing  agent,  both  in  the  industries 
and  in  the  laboratory.  Its  oxidizing  action  is  easily  understood  when 
it  is  remembered  that  manganese  can  play  the  part  either  of  an  acid- 
forming  element  of  high  valence  (7)  or  of  a  base-forming  element  of 
lower  valence  (2  or  4).  The  decomposition  of  the  permanganate  may 
take  place  in  two  different  ways,  depending  upon  whether  the  reaction 
occurs  in  an  acid  solution  or  in  one  that  is  neutral  or  basic. 

1.  Oxidation  in  acid  solution.  When  the  permanganate  is  brought 
into  a  solution  of  an  acid,  both  the  potassium  and  the  manganese 
tend  to  form  salts  with  the  given  acid.  This  brings  about  a  complete 
rearrangement  of  the  constituents  of  the  permanganate,  which  results 
in  the  liberation  of  oxygen  : 

2  KMn04  +  3  H2SO4  =  K2SO4  +  2  MnSO4  +  3  H2O  +  5  O     (1) 


MANGANESE  AND  CHROMIUM  519 

This  action  is  not  very  noticeable  unless  some  reducing  agent  is 
present  to  take  up  the  oxygen,  under  which  conditions  it  is  very 
rapid,  even  at  ordinary  temperatures.  Since  the  permanganate  solution 
is  intensely  colored,  while  the  products  of  oxidation  are,  as  a  rule, 
almost  colorless,  the  completion  of  the  oxidation  can  be  determined 
with  great  accuracy.  This  combination  of  properties  makes  potassium 
permanganate  of  the  greatest  service  in  chemical  analysis.  The  fol- 
lowing equations  illustrate  the  oxidizing  power  of  the  permanganate : 

Oxalic  acid :  C2H2O4  +  O  =  2  CO2  +  H2O  (2) 

Ferrous  sulfate :  2  FeSO4  +  H2SO4  +  O  =  Fe2(SO4)3  +  H2O        (3) 
Hydrogen  chloride :  2  HC1  +  O  =  C12  +  H2O  (4) 

The  complete  equations  can  be  obtained  by  combining  these  with 
(1)  given  above.  In  the  case  of  ferrous  sulfate  it  is  as  follows : 

2  KMnO4  +  10  FeSO4  +  8  H2SO4 

=  5  Fe2(S04)3  +  K2S04  +  2  MnSO4  +  8  H2O 

2.  Oxidation  in  neutral  or  basic  solution.  In  neutral  or  basic  solu- 
tion the  tendency  is  for  the  manganese  to  become  tetravalent,  forming 
the  compound  Mn(OH)4.  The  reaction,  which  takes  place  with  notice- 
able rapidity  only  in  the  presence  of  a  reducing  agent,  is  expressed 
in  the  following  equation: 

2  KMnO4  +  5  H2O  =  2  Mn(OH)4  +  2  KOH  +  3  O 
If  alcohol  is  added,  it  is  oxidized  to  acetic  acid  : 

C2H5OH  +  20  =  C2H402  +  H20 
The  combined  equation  is  as  follows : 

4  KMnO4  + 10  H2O  +  3  C2H5OH 

-  4  Mn(OH)4  +  4  KOH  +  3  C2H4O2  +  3  H2O 

CHROMIUM 

Chromium  was  discovered  by  the  French  chemist  Vauquelin  in 
1797,  during  an  investigation  of  the  rather  rare  mineral  crocoite 
(PbCrO4).  Its  chief  occurrence  in  nature  is  in  the  form  of  chromite, 
or  chrome  iron  ore  (Fe(CrO2)2),  but  in  traces  it  is  rather  widely  dis- 
tributed in  minerals,  many  of  which,  such  as  the  emerald,  appear  to  owe 
their  green  color  to  its  compounds.  The  commercial  supply  of  chrome 
ore  comes  from  Rhodesia,  New  Caledonia,  and  Greece,  with  smaller 
quantities  from  California. 


520  GENERAL  CHEMISTRY 

Preparation  and  properties.  The  pure  metal  is  best  prepared  by  the 
Goldschmidt  process,  and  a  somewhat  less  pure  product  by  careful 
reduction  of  the  oxide  by  carbon  in  an  electric  furnace,  any  excess  of 
carbon  being  avoided.  It  is  a  highly  crystalline,  brilliant,  silvery  metal, 
very  hard  and  brittle.  Specimens  containing  a  little  carbon  are  very 
much  harder  than  the  pure  metal.  Its  density  is  6.50,  its  melting  point 
1489°,  and  its  boiling  point  about  2200°. 

The  element  is  unoxidized  in  air  at  all  ordinary  temperatures,  but 
when  it  is  finely  powdered  and  sufficiently  heated,  it  burns  with  great 
brilliancy.  It  displaces  hydrogen  from  dilute  acids,  forming  chromous 
salts,  but  it  is  not  attacked  by  oxidizing  acids  like  nitric  acid,  assum- 
ing instead  a  passive  condition  similar  to  that  of  iron. 

Uses.  The  metal  is  quite  extensively  used  in  the  steel  industry, 
since  its  alloys  with  iron  are  very  hard  and  well  adapted  to  special 
uses.  For  such  purposes  the  metal  is  prepared  in  the  form  of  ferro- 
chromium  by  the  reduction  of  chrome  iron  ore  with  carbon,  the  product 
containing  from  60  to  70  per  cent  of  chromium  and  from  1  to  8  per 
cent  of  carbon.  Its  alloys  with  nickel,  cobalt,  and  copper  are  also  very 
hard  and  strong,  the  one  consisting  of  25  per  cent  chromium  and  75  per 
cent  cobalt  being  especially  well  adapted  to  the  manufacture  of  cutlery. 

Compounds  of  chromium.  Like  manganese,  chromium  can  exist  in  a 
number  of  stages  of  oxidation,  its  chief  compounds  being  derived  from 
oxides  of  the  formulas  CrO,  Cr2O8,  and  CrOg.  The  first  of  these  is 
always  basic ;  the  second,  like  the  corresponding  oxide  of  aluminium, 
is  predominantly  basic,  but  to  a  limited  extent  plays  the  part  of  an  acid 
anhydride  ;  the  third  is  always  acid  in  character.  Nearly  all  the  com- 
pounds of  chromium  are  highly  colored  in  shades  of  blue,  green,  violet, 
red,  and  yellow  —  a  fact  which  originally  suggested  the  name  of  the 
element,  from  the  Greek  word  meaning  "  color." 

.  Chromous  compounds.  Chromous  salts,  corresponding  to  the  oxide 
CrO,  are  most  easily  prepared  by  dissolving  chromium  in  the  appro- 
priate acid,  the  hydrogen  evolved  in  the  reaction  preventing  the  salt 
from  oxidizing.  Solutions  containing  them  can  also  be  prepared  by 
the  reduction  of  chromic  salts  by  means  of  such  metals  as  zinc : 

2  CrCl8  +  Zn  =  2  CrCl2  +  ZnCla 

If  sodium  acetate  is  added  to  the  solution  so  obtained,  the  difficultly 
soluble  chromous  acetate  (Cr(C2H3O2)2  •  H2O)  is  obtained  in  the  form 
of  deep-red  crystals,  and  this  is  the  best-known  chromous  salt  in  the 


MANGANESE  AND  CHROMIUM  521 

solid  state. r  The  corresponding  hydroxide  (Cr(OH)2)  is  obtained  as 
a  yellow  precipitate,  which  oxidizes  with  such  ease  that  it  slowly 
decomposes  water,  with  evolution  of  hydrogen : 

2  Cr(OH)2  +  2  H20  =  2  Cr(OH)3  +  H2 
Chromous  salts  act  in  the  same  general  way  : 

6  CrCl2  4-  6  H20  -  4  CrCl3  +  2  Cr(OH)3  +  3  H2 

Chromic  compounds.  The  compounds  derived  from  chromic  oxide 
(Cr2O3)  are  the  stable  compounds  of  chromium  as  a  base-forming 
element.  They  are  analogous  to  the  corresponding  salts  of  aluminium 
and  ferric  iron,  and,  like  these,  are  quite  extensively  hydrolyzed  in 
solution.  Chromic  salts  of  most  of  the  familiar  acids  have  been  de- 
scribed. Most  of  these  exhibit  a  peculiarity  not  encountered  in  other 
salts  except  in  the  case  of  some  of  the  rarer  elements ;  they  exist  in 
two  modifications,  which  have  the  same  formulas  but  are  very  different 
in  properties,  as  will  be  explained  in  connection  with  the  chloride. 

Chromic  oxide  (Cr203).  This  oxide  can  be  obtained  by  the  various 
methods  available  in  the  preparation  of  oxides ;  it  is  a  green  powder, 
the  exact  shade  of  which  depends  upon  its  physical  condition.  In  the 
crystalline  state  it  is  very  bright  green  in  color,  but  otherwise  it  bears 
a  strong  resemblance  to  corundum.  For  use  as  a  pigment  it  is  usually 
prepared  by  heating  sodium  dichromate  with  a  suitable  reducing  agent, 
sulfur  serving  very  well : 

Na2Cr20?  +  S  =  Na2SO4  +  O2O3 

The  hydroxide  (Cr(OH)3),  as  prepared  by  precipitation,  is  a  rather  dark 
green,  amorphous  solid,  which  gives  a  number  of  dehydration  products. 
Chromic  chloride  (CrCl3).  This  compound  will  serve  very  well  to 
illustrate  some  of  the  peculiarities  of  chromic  salts  in  general.  It  can 
be  obtained  in  the  form  of  beautiful  violet-colored,  pearly  scales  by 
preparing  it  in  the  absence  of  water,  as  by  heating  a  mixture  of 
chromic  oxide  and  carbon  in  a  current  of  chlorine : 

Cr203  +  3  C  +  3  C12  =  2  CrCl3  +  3  CO 

It  is  apparently  insoluble  in  water,  but  it  dissolves  upon  long  stand- 
ing or,  much  more  rapidly,  through  the  catalytic  action  of  a  trace 
of  a  chromous  compound,  forming  a  green  solution.  Under  favorable 
conditions  green  crystals  deposit  from  this  solution,  having  the  com- 
position expressed  in  the  formula  CrCl  •  6  HO.  If  silver  nitrate  is 


522  GENERAL  CHEMISTRY 

added  to  this  solution,  only  two  thirds  of  the  chlorine  is  precipitated, 
which  indicates  that  one  of  the  chlorine  atoms  is  in  a  different  condi- 
tion from  the  other  two  and  is  not  an  ion.  Upon  long  standing  the 
solution  turns  violet  in  color  and  deposits  crystals  of  a  gray-blue  tint. 
These  also  have  the  formula  CrCl3  •  6  H2O,  and  their  solution  acts 
normally  with  silver  nitrate. 

Many  other  chromic  salts  exist  in  two  modifications  analogous  to 
these  two  soluble  salts.  Long  standing  and  low  temperature  favor 
the  formation  of  the  violet  form,  while  rapid  formation  and  hot  solu- 
tions produce  the  green  variety.  In  many  instances  there  is  doubtless 
hydrolysis,  which  results  in  the  formation  of  green  basic  salts  ;  but  that 
there  are  complexes  of  other  kinds  is  shown  by  the  fact  that  the  green 
modifications  rarely  act  normally  with  precipitating  reagents. 

Other  chromic  salts.  Of  the  other  chromic  salts  the  sulfate 
(Cr2(SO4)3)  is  the  best  known.  The  violet  form  has  the  composition 
Cr2(SO4)3'15H2O.  A  number  of  green  modifications  with  varying 
compositions  are  known.  When  crystallized  in  the  presence  of  potas- 
sium or  ammonium  sulfate,  chromic  sulfate  forms  an  alum,  which 
yields  large,  ruby-colored  octahedra  of  great  perfection  of  form.  Po- 
tassium chrome  alum  (KCr(SO4)2'  12H0O)  has  extensive  use  in  the 
leather  industry.  In  solution  it  yields  chromic  hydroxide  by  hydrolysis, 
and  this  compound  acts  upon  animal  skins  in  much  the  same  way  as 
tanbark  does,  the  tanning  process  being  accomplished  very  much  more 
rapidly  than  when  bark  is  used. 

Like  aluminium  and  ferric  iron,  chromium  forms  no  chromic  car- 
bonate, nor  does  it  form  a  sulfide  by  precipitation  methods.  Reac- 
tions which  would  normally  lead  to  the  formation  of  these  compounds 
produce  the  hydroxide  instead,  as  is  true  with  most  of  the  trivalent 
metals. 

Chromites.  Chromium  hydroxide,  like  aluminium  hydroxide,  is 
soluble  in  excess  of  the  alkalies,  with  the  formation  of  chromites. 
These  are  derived  from  chromium  hydroxide,  or  its  partial  anhydride, 
acting  as  acid  :  CrQ  .  QH  + 


Soluble  salts  of  this  kind  are  readily  hydrolyzed,  yielding  chromic 
hydroxide  in  colloidal  form.  When  the  solution  is  boiled,  the  hydrox- 
ide is  coagulated  as  a  green  precipitate.  A  number  of  salts  of  the 
acid  HCrO2  are  found  in  nature,  chief  among  which  is  the  ferrous 
salt,  chrome  iron  ore  (Fe(CrO  )  ). 


MANGANESE  AND  CHROMIUM  523 

Chromic  acid  (H2CrO  J ;  the  chromates.  When  any  of  the  compounds 
of  chromium  so  far  mentioned  are  heated  with  an  alkali  (or  alkaline 
carbonate),  oxygen  is  absorbed  from  the  air,  and  the  chromium  be- 
comes hexavalent,  forming  salts  of  chromic  acid  (H2CrO4).  With  iron 
chromite  the  reaction  is  represented  in  the  equation 

4  Fe(Cr02)2  +  16  KOH  +  7  O2  =  8  K2CrO4  +  2  FeaOt  +  8  H2O 

In  formula  and  crystalline  form  these  salts  are  analogous  to  the  sul- 
fates  and  manganates.  Most  of  them  are  of  some  shade  of  yellow, 
unless  the  base  present  contributes  a  color  of  its  own  to  the  saltt 
Of  the  chromates  the  potassium  salt  (K2CrO4)  and  the  sodium  salt 
(Na2CrO4  •  10  H2O)  are  the  best-known  soluble  representatives. '  The 
insoluble  lead  chromate  (PbCrO4)  is  the  pigment  known  as  chrome 
yellow.  The  insoluble  barium  salt  (BaCrO4)  is  of  a  lighter  shade  of 
yellow  and  is  also  used  as  a  pigment. 

The  free  chromic  acid  cannot  be  prepared  in  pure  condition,  owing 
to  its  tendency  to  lose  water,  with  the  formation  of  various  condensed 
soluble  acids,  the  chief  of  which  is  dichromic  acid  (H2Cr2O?).  These 
also  lose  water  readily,  finally  forming  the  trioxide  (CrO3)  which  crys- 
tallizes in  deep  red  needles,  very  soluble  in  water  but  much  less  so 
in  sulfuric  acid.  Consequently,  when  concentrated  sulfuric  acid  is 
added  to  a  chromate,  it  is  this  anhydride  which  is  obtained  and  not 
one  of  the  chromic  acids  : 

K2Cr04  +  H2S04  =  K2S04  +  H2O  +  CrO3 

Dichromic  acid  (H2Cr207);  the  dichromates.  Unlike  sulfuric  acid, 
chromic  acid  forms  no  acid  salts.  Eeactions  which  might  be  expected 
to  produce  such  salts  lead  instead  to  salts  of  various  condensed  acids, 
the  chief  of  which  is  dichromic  acid  (H2Cr2O?).  Thus,  when  the  cal- 
culated quantity  of  sulfuric  acid  acts  upon  a  solution  of  potassium 
chromate,  the  dichromate  crystallizes  from  the  solution  in  the  form  of 
orange-red,  triclinic  crystals  of  the  composition  K2Cr2O7,  the  salt  being 
analogous  to  potassium  pyrosulfate  (K2S2O7)  : 

2  K2Cr04  +  H2S04  =  K2O2O7  +  K2SO4  +  H2O 

This  salt  is  very  soluble  in  hot  water  but  only  moderately  soluble  in 
cold.  On  account  of  its  moderate  solubility  and  easy  purification  by 
crystallization,  this  salt  has  long  been  used  for  all  purposes  to  which 
a  soluble  derivative  of  chromic  acid  is  suited.  The  sodium  salt 


524  GENERAL  CHEMISTRY 

(Na2Cr2O7  •  2  H2O)  is  very  soluble  and  somewhat  deliquescent,  but 
the  difficulties  attending  its  preparation  have  been  overcome,  and  for 
commercial  purposes  the  sodium  compound  is  now  replacing  the 
potassium  salt. 

When  a  solution  of  a  dichromate  is  treated  with  an  excess  of  alkali, 
the  normal  chromate  is  obtained,  just  as  would  be  the  case  with  an 
acid  salt.  Excess  of  acid  converts  the  chromate  once  more  into  the 
dichromate,  so  that  the  condition  in  solution  is  one  of  equilibrium,  as 
expressed  in  the  equation 

2  K2Cr04  +  H20  +=*  K2Cr2O7  +  2  KOH 

The  anions  of  both  chromate  and  dichromate  are  present  in  the  solu- 
tion, and  when  various  salts  are  added,  either  a  chromate  or  a  dichro- 
mate may  precipitate,  depending  upon  which  is  the  less  soluble.  For 
example,  normal  lead  chromate  is  precipitated  when  a  solution  of  a 
dichromate  is  treated  with  a  lead  salt : 

2Pb(C2H802)2  +  K20207  +  H20   . 

=  2  PbCr04  +  2  KC2H802  +  2  HC2H3O2 

Oxidizing  action  of  the  chromates.  Owing  to  the  fact  that  chromium, 
like  manganese,  can  readily  diminish  in  valence  and  play  the  part  of 
a  base-forming  element,  both  the  chromates  and  the  dichromates  are 
good  oxidizing  agents.  The  oxidation  reaction  in  the  presence  of  an 
acid  —  for  example,  sulfuric  acid  —  is  represented  in  the  equations 

2  K2Cr04  +  5  H2S04  -  2  K2SO4  +  Cra(SO4),  +  5  H2O  +  3  O 
K2CrA  +  4  H2S04  =  K2S°4  +  Cra(S04)8  +  4  H20  +  3  O 

These  reactions  take  place  to  a  noticeable  extent  only  when  some  re- 
ducing agent  is  present  to  take  up  the  oxygen  represented  as  being 
liberated.  For  example,  hydrogen  chloride  is  oxidized  to  water  and 
chlorine,  alcohol  to  acetic  acid,  and  ferrous  salts  to  ferric.  The  reac- 
tion in  the  case  of  hydrogen  chloride  (p.  249)  is  represented  in  the 
equation 

K2Cr207  +  4  H2S04  +  6  HC1  =  K2SO4  +  Cr2(SO4)3  +  7  H2O  +  3  C12 

In  neutral  or  basic  solution  both  the  potassium  and  the  chromium 
are  converted  into  hydroxides  when  the  chromate  acts  as  an  oxidizing 
reagent.  With  the  dichromate  the  equation  is  as  follows : 

K2Cr2Or  +  4  H20  =  2  KOH  +  2  Cr(OH)3  +  3  O 


CHAPTER  XXXVIII 

THE  VANADIUM  AND  MOLYBDENUM  FAMILIES 

Metallo-acid  elements.  In  several  places  in  the  preceding  pages 
attention  has  been  directed  to  the  fact  that  the  base-forming  elements 
differ  very  markedly  from  the  acid-forming  elements  in  their  proper- 
ties, and  to  the  further  fact  that  some  of  the  metallic  elements  in 
higher  stages  of  oxidation  form  acids.  This  is  particularly  true  in 
the  case  of  tin,  chromium,  and  manganese,  which,  as  elements,  are 
preeminently  metals,  and  in  lower  stages  of  oxidation  act  as  bases, 
yet  form  well-defined  acids. 

Among  the  less  familiar  elements  there  are  a  number  which  act  in 
this  same  way.  The  elementary  substances  have  all  of  the  properties 
usually  associated  with  the  term  metal,  but  in  their  chemical  conduct 
they  play  the  part  of  acid-forming  elements.  In  lower  stages  of  oxi- 
dation some  of  them  are  like  chromium  and  manganese  in  that  they 
act  as  bases  and  form  salts  with  the  common  acids,  but  they  usually 
tend  to  pass  to  a  higher  state  of  oxidation,  in  which  they  form  a  series 
of  acids.  Some  of  them  do  not  act  as  bases  at  all,  but  are  exclusively 
acid-forming  in  character.  The  term  metallo-acid  is  often  applied  to 
elements  of  this  kind,  and  to  some  extent  it  is  applicable  to  chro- 
mium and  manganese.  It  is  more  frequently  employed  to  designate 
the  family  in  the  fifth  group,  consisting  of  vanadium,  columbium,  and 
tantalum;  to  the  elements  following  chromium  in  the  sixth  group, 
namely,  molybdenum,  tungsten,  and  uranium ;  and  to  the  elements  in 
the  eighth  group,  known  collectively  as  the  platinum  metals.  The  first 
two  of  these  groups  will  be  considered  in  the  present  chapter,  the 
platinum  metals  being  reserved  for  the  next. 

General.  The  three  elements,  vanadium,  niobium,  and  tantalum, 
named  in  honor  of  mythological  personages,  constitute  a  family  in  the 
group  with  phosphorus.  In  English-speaking  countries  the  second 
one  is  more  commonly  named  columbium,  in  honor  of  America,  since 
it  was  discovered  in  a  mineral  from  Connecticut.  It  was  rediscovered 
as  an  impurity  in  tantalum  and  renamed  niobium  (Niobe  being  the 
daughter  of  Tantalus),  and  both  names  are  in  use. 

525 


526  GENERAL  CHEMISTRY 

Vanadium.  Vanadium  was  first  described  as  a  new  element  by 
Sef strom  in  1830,  and  named  in  honor  of  Vanadis,  the  Scandinavian 
goddess  of  fortune.  It  is  much  the  most  abundant  of  the  three  and 
is  widely  distributed  in  nature,  but  it  is  nowhere  concentrated  in  very 
rich  deposits.  Vanadinite  (Pb3(VO4)2  •  PbCl2)  is  the  most  abundant 
mineral,  occurring  in  the  form  of  reddish,  crystalline  crusts  upon 
many  other  minerals.  At  present  much  of  the  vanadium  of  commerce 
comes  from  the  ashes  of  certain  Peruvian  coals,  which  run  as  high  as 
48  per  cent  vanadic  oxide,  and  from  some  South  American  ores,  which 
are  chiefly  sulfides. 

Preparation.  Metallic  vanadium  is  Tery  difficult  to  prepare,  owing 
to  the  fact  that  it  is  reducible  only  at  a  very  high  temperature,  and 
under  these  conditions  it  combines  with  most  reducing  agents  and 
with  nitrogen.  It  is  most  successfully  prepared  either  by  reducing  its 
oxide  with  mixed  metal  (p.  452)  or  by  a  method  devised 
by  von  Bolton  and  applicable  to  the  other  metals  of 
this  group  and  to  tungsten. 

Method  of  von  Bolton.  This  method  depends  upon  the  fact 
that  the  lower  oxides  of  these  metals  are  conductors  of  the 
electric  current,  but  that,  when  highly  heated  by  the  current, 
they  decompose,  yielding  the  metal  and  free  oxygen.  The 
pentoxide  (V2O5)  is  mixed  with  paraffin  and  carbon,  molded 
into  the  form  of  a  wire  or  rod,  and  baked  at  a  high  tem- 
perature, by  which  process  reduction  to  the  trioxide  (V2O3)  is 
effected.  The  fragile  rod  so  obtained  (Fig.  154,^1)  is  sealed 
in  a  glass  bulb  D  capable  of  being  exhausted  by  an  air  pump 
attached  at  E.  Electrical  connection  is  provided  by  the  wires 
B,  B,  which  enter  through  the  glass  rod  C  as  in  an  incandes- 
cent lamp.  A  moderately  strong  current  is  passed  through  the  rod,  while  the 
pump  removes  the  oxygen  as  it  escapes.  In  this  way  a  metallic  wire  is  formed. 

Properties  and  chemical  conduct.  The  pure  metal  is  a  brilliant,  sil- 
very substance,  crystallizing  in  the  hexagonal  system  like  phosphorus. 
It  is  very  hard  and  somewhat  brittle,  and  has  a  density  of  about  6.02. 
It  melts  at  1620°.  By  reducing  a  mixture  of  the  oxides  of  iron  and 
vanadium  an  alloy  called  ferrovanadium  is  obtained.  This  alloy  is 
used  in  the  manufacture  of  vanadium  steel,  which  is  very  tough  and 
strong.  This  application  constitutes  the  chief  use  for  vanadium, 
though  some  of  its  compounds  are  employed  as  catalytic  agents,  as  in 
the  manufacture  of  certain  dyes  and  of  sulfuric  acid.  The  metal  is 
rather  easily  oxidized  and  is  soluble  in  concentrated  acids,  forming 
metallic  salts. 


THE  VANADIUM  AND- MOLYBDENUM  FAMILIES      527 

Compounds  of  vanadium.  Compounds  of  vanadium,  like  those  of 
chromium,  are  obtained  by  heating  the  ore  with  sodium  carbonate  and 
an  oxidizing  agent  (p.  523),  sodium  vanadate  (Na3VO4)  being  formed 
in  the  reaction.  When  a  solution  of  this  salt  is  treated  with  an 
excess  of  ammonium  chloride,  the  sparingly  soluble  ammonium 
metavanadate  (NH4VO3)  is  precipitated: 

Na8VO4  +  3  NH4C1  =  NH4VO3  +  3  NaCl  +  2  NH8  +  H2O 

When  this  salt  is  heated,  the  pentoxide  (V2O6)  is  obtained,  and  from 
this  other  compounds  can  be  made.  Many  of  these  are  highly  colored 
in  shades  of  violet,  blue,  green,  red,  and  yellow.  The  vanadic  acids 
correspond  to  those  of  phosphorus,  the  simplest  ones  having  the  for- 
mulas HVO8,  H3VO4,  and  H4V2O7,  but  many  condensed  acids  are 
also  known.  Vanadium  forms  five  oxides,  corresponding  to  those  of 
nitrogen.  As  a  divalent  metal  it  forms  a  deep  blue  vitriol  of  the  for- 
mula VSO4  •  7  H2O  and  a  double  ammonium  sulfate  of  the  formula 
(NH4)2SO4  •  VSO4  •  6  H2O.  As  a  trivalent  metal  it  forms  alums,  which 
range  in  color  from  the  greenish-violet  potassium  alum  to  the  ruby- 
red  caesium  alum.  The  blue,  liquid  tetrachloride  is  formed  by  the 
action  of  hydrochloric  acid  upon  the  pentoxide.  It  solidifies  only  at 
the  temperature  of  liquid  air.  As  usually  prepared,  the  pentoxide  is 
a  cherry-red,  crystalline  body.  The  great  majority  of  the  derivatives 
of  vanadium  are  complexes,  sometimes  of  very  complicated  formulas. 

Columbium  and  tantalum.  These  two  elements  are» usually,  though 
not  always,  found  together  in  nature  in  the  form  of  columbates  and 
tantalates.  Minerals  containing  them  are  found  in  many  localities, 
especially  in  the  United  States,  but  they  are  very  local  and  never 
occur  in  large  deposits.  Columbite,  the  mineral  in  which  they  were 
first  discovered,  is  essentially  an  iron  and  manganese  salt  of  the  mixed 
acids  HCbO3  and  HTaO3,  and  is  much  richer  in  columbium  than  in 
tantalum.  Tantalite,  chiefly  of  the  composition  Fe(TaO3)2,  always 
carries  some  columbium.  Samarskite  is  a  wonderfully  complex  mineral, 
the  acids  in  which  are  chiefly  columbic  and  tantalic,  while  almost  all 
of  the  rare  earths,  as  well  as  many  other  oxides,  are  present  as  bases. 

Preparation  and  properties.  The  metals  are  obtained  by  the  methods 
described  in  connection  with  vanadium,  columbium  being  much  the 
more  difficult  to  prepare  in  a  pure  state.  It  was  first  so  obtained  in 
1905,  by  von  Bolton,  who  found  it  to  be  a  steel-gray  metal  of  brilliant 
luster,  about  as  hard  as  wrought  iron  and  rather  malleable  and 


528  GENERAL   CHEMISTRY 

ductile.  Heated  in  the  air,  it  is  slowly  oxidized.  It  is  very  little  acted 
upon  by  the  common  acids,  but  is  attacked  by  fused  alkalies,  and  it 
combines  directly  with  hydrogen  and  nitrogen  at  higher  temperatures. 
Its  density  is  12.75  and  its  melting  point  1950°. 

Tantalum  resembles  platinum  in  appearance  and  is  about  as  hard  as 
wrought  iron.  It  is  more  ductile  than  columbium,  so  that  von  Bolton 
was  able  to  draw  wires  only  0.03  mm.  in  diameter.  Its  density  is  16.6 
and  its  ^melting  point  is  very  high,  being  estimated  by  Burgess  at 
2910°.  It  is  inactive  toward  alkaline  solutions  and  toward  acids 
other  than  hydrofluoric.  At  a  low  red  heat  it  burns  in  the  air. 

The  ductility  of  tantalum  and  its  very  high  melting  point  make  it 
well  adapted  to  serve  as  a  metallic  filament  in  incandescent  lamps. 
Since  it  is  a  good  conductor,  the  wire  must  be  very  fine  and  much 
longer  than  an  ordinary  carbon  filament,  in  order  that  the  resistance 
may  be  great  enough  to  keep  the  current  at  the  desired  value  for  in- 
candescence. The  filament  is  therefore  rather  fragile  and  tends  to  bend 
under  its  own  weight.  It  has  been  almost  entirely  replaced  by  tungsten. 

Compounds  of  the  elements.  The  compounds  of  columbium  and 
tantalum  are  for  the  most  part  very  complex.  The  simplest  ones  are 
the  oxides  and  the  various  halogen  derivatives,  particularly  the  fluo- 
acids  and  their  salts.  These  acids  have  different  formulas,  namely, 
H2CbF5O  and  H2TaF7,  and  the  different  solubility  of  salts  of  these 
two  acids  affords  the  best  means  of  separating  the  two  elements.  Their 
simple  oxygen  acids  are  like  those  of  vanadium  in  formula,  but  for  the 
most  part  their  salts  are  derived  from  a  variety  of  condensed  acids, 
such  as  those  represented  by  the  formulas  HgCb6O19  and  HgTa6O19. 

Molybdenum,  tungsten,  and  uranium.  The  elements  which  follow 
chromium — namely,  molybdenum,  tungsten,  and  uranium  —  are  rather 
more  abundant  in  nature  than  those  of  the  preceding  group.  As  ele- 
ments they  are  metals  of  high  melting  point  and  density.  Molybdenum 
and  tungsten  are  almost  entirely  acid-forming  in  chemical  conduct, 
while  uranium  combines  both  acid-forming  and  base-forming  qualities. 
Like  chromium,  they  are  usually  hexavalent,  and  many  of  their  com- 
pounds have  formulas  similar  to  those  of  chromium. 

Molybdenum.  The  Greek  word  from  which  molybdenum  is  derived 
was  applied  in  early  writings  to  many  substances  having  a  superficial 
resemblance  to  lead  —  among  others,  to  a  mineral  resembling  graphite, 
which  is  now  called  molybdenite.  In  1778  Scheele  showed  that  from 
this  mineral  a  new  oxide  could  be  obtained,  which  he  called  molybdic 


THE  VANADIUM  AND  MOLYBDENUM  FAMILIES      529 

acid.  The  element  occurs  chiefly  in  molybdenite  (MoS2),  in  wulfenite 
(PbMoO4),  and  in  molybdic  ocher,  a  hydrated  molybdate  of  iron.  Its 
ores  are  nowhere  very  abundant  and  usually  require  much  concentra- 
tion before  they  are  suitable  for  metallurgical  purposes. 

Preparation  and  properties.  Molybdic  oxide  (MoO3)  can  be  reduced 
by  hydrogen  to  the  state  of  a  fine  metallic  powder,  but  the  compact 
form  of  the  metal  is  obtained  by  reduction  with  aluminium  or,  better, 
with  mixed  metal.  It  is  a  moderately  heavy,  silvery  metal,  of  density 
8.6,  melting  only  at  a  white  heat.  It  is  too  volatile  to  serve  as  a  fila- 
ment for  lamps,  but  in  the  form  of  a  fine  wire  it  is  used  as  a  support 
for  tungsten  filaments.  To  some  extent  it  is  used  as  an  alloy  metal  in 
special  varieties  of  steel. 

Compounds.  Molybdenum  forms  a  great  variety  of  oxides,  ranging 
from  the  monoxide  (MoO)  to  the  trioxide  (MoO3).  The  latter  is  a 
nearly  white,  crystalline  compound,  which  is  rather  easily  volatile  and 
is  frequently  called  molybdic  acid.  The  simple  molybdates  derived  from 
it  have  formulas  like  the  chromates, —  for  example,  the  potassium  salt 
(K2MoO4), —  but  usually  the  salts  are  derived  from  condensed  acids, 
ammonium  molybdate  having  the  formula  (NH4)6Mo7O24  •  4  H2O. 

Complex  acids.  When  molybdic  acid  is  brought  into  contact  with 
soluble  salts  of  phosphoric  acid,  especially  •  in  the  presence  of  nitric 
acid,  various  complex  salts  are  deposited,  the  formulas  of  which  are 
very  complicated.  They  are  usually  expressed  by  writing  the  formulas 
of  the  anhydrides  of  both  acids  and  base,  as  shown  in  the  formula  for 
ammonium  phosphomolybdate,  3  (NH4)2O  •  P2O5  •  22  MoO3  •  12  H2O. 
Similar  salts  are  formed,  in  which  the  P2O5  and  the  MoO3  are  replaced 
by  a  number  of  different  anhydrides,  such  as  As2O5, 12Og,  Ta2O5,  V2O5, 
WO3,  UO3,  and  many  others.  The  ratios  are  sometimes  very  compli- 
cated, and  nothing  is  known  about  the  way  in  which  these  compounds 
are  made  up ;  yet  they  are  often  very  beautifully  crystallized  and 
appear  to  be  perfectly  definite  chemical  compounds.  The  ammonium 
salt,  whose  formula  is  given  above,  is  a  yellow,  crystalline  precipitate, 
insoluble  in  nitric  acid,  and  advantage  is  taken  of  its"  formation  in  the 
separation  and  estimation  of  phosphoric  acid.  This  constitutes  the 
most  important  use  of  molybdenum  compounds. 

Tungsten.  The  element  tungsten  was  discovered  by  Scheele  as  a 
constituent  of  a  mineral  now  called  scheelite  (CaWO4).  This  mineral 
and  wolframite  (FeWO4)  are  not  very  rare  in  nature,  and  considerable 
quantities  of  the  ore  are  produced  annually.  Its  compounds  are  easily 


530  GENERAL  CHEMISTRY 

extracted  from  the  ore.  It  is  not  difficult  to  reduce  the  oxide  to  the 
metallic  state,  but  a  powder  is  obtained  which  it  is  almost  impossible 
to  melt  to  a  compact  ingot.  It  is  best  prepared  like  vanadium,  or 
the  powdered  metal  can  be  dissolved  in  mercury  and  cadmium,  the 
amalgam  fashioned  into  wire*  and  the  other  metals  expelled  by  elec- 
trical heating,  leaving  the  pure,  coherent  tungsten.  A  similar  result 
can  be  attained  by  pressing  the  powder  into  the  form  of  a  wire  and 
rolling  or  hammering  it  at  a  high  temperature,  maintained  by  electrical 
heating.  It  is  a  very  hard  metal,  as  ordinarily  produced,  but  by  the 
proper  mechanical  treatment  it  can  be  obtained  in  a  soft  and  malle- 
able state  resembling  platinum.  It  has  a  very  high  melting  point, 
estimated  at  3000°,  is  a  good  conductor,  and  does  not  oxidize  very 
readily  in  the  air.  It  is  of  much  value  for  spark  points  for  automobiles, 
electrical  contacts  for  large  currents,  and  filaments  in  incandescent 
lamps.  It  is  proposed  to  use  it  in  the  manufacture  of  bullets,  since  its 
high  density  (which  is  from  17  to  19)  would  be  of  great  advantage. 
Tungsten  is  used  as  an  important  alloy  metal  for  steels,  particularly 
for  armor  plate  and  for  tools.  Tungsten  steel  holds  its  temper  at  a 
red  heat,  so  that  with  tools  of  this  alloy  a  lathe  can  be 
driven  at  a  very  high  speed  without  injury  to  the  tool. 

Tungsten  lamp.  The  efficiency  of  any  material  as  a  filament  in 
an  incandescent  lamp  depends  primarily  upon  its  conductivity  and 
its  melting  point.    Other  conditions  being  the  same,  most  sub- 
stances are  equally  incandescent  at  a  given  temperature,  and  the 
higher  the  temperature  the  more  brilliant  the  light  they  give. 
Tungsten    not   only  has   the    highest    melting  point  of   all  the 
metals  but  it  has  practically  no  vapor  pressure  below  its  melting 
point.    It  is  a  good  conductor,  so  that  to  secure  sufficient  resist- 
FIG.  155        ance  it  is  necessary  to  use  a  wire  that  is  very  fine  as  well  as 
rather  long.     This  necessitates  looping  the  wire,  as    shown   in 
Fig.  155,  and  supporting  it  at  frequent  intervals  to  prevent  sagging.  While  carbon 
does  not  melt  at  any  attainable  temperature,  it  volatilizes  so  easily  that  it  is  not 
practicable  to  heat  a  carbon  filament  above  a  bright  yellow  glow.    The  relative 
efficiency  of  some  common  types  of  lamps  is  approximately  as  follows : 

Ordinary  carbon  filament 3.25  watts  per  candle  power 

Graphitized  carbon  filament  (Gem)     .     .  2.50  watts  per  candle  power 

Tantalum  filament 2.00  watts  per  candle  power 

Tungsten  filament 1.25  watts  per  candle  power 

Compounds  of  tungsten.  In  a  general  way  the  compounds  of  tungsten 
resemble  those  of  molybdenum.  A  few  form  valuable  pigments ;  so- 
dium tungstate  (Na2WO4)  is  used  as  a  fireproof  coating  for  cloth,  and 


THE  VANADIUM  AND  MOLYBDENUM  FAMILIES      531 

also  as  a  mordant  in  the  dyeing  industry.  Phosphotungstic  acid 
(analogous  to  phosphomolybdic  acid)  is  used  as  a  reagent  in  the 
detection  of  certain  alkaloids. 

Uranium.  Uranium  was  discovered  by  Klaproth,  in  1789,  in  the 
mineral  known  as  pitchblende ;  it  was  named  in  honor  of  the  planet 
Uranus,  which  had  been  discovered  only  a  short  time  before.  It  is 
found  in  quite  a  number  of  complex  minerals  associated  with  lead,  vana- 
dium, thorium,  and  the  rare  earths.  Of  these,  uraninite,  or  pitchblende 
(U8Og),  is  the  most  common,  the  ore  sometimes  carrying  from  75  to  85 
per  cent  uranium.  A  ton  of  pitchblende  is  valued  at  about  ten  thousand 
dollars ;  Joachimstal,  in  Bohemia,  leads  in  the  production  of  this  mineral. 

Properties  and  conduct.  Uranium  is  a  somewhat  malleable  metal  re- 
sembling nickel  in  appearance.  It  oxidizes  slowly  in  the  air  and  is 
soluble  in  the  ordinary  acids,  forming  salts.  Its  density  is  18.7  and  its 
melting  point  about  1500°.  As  an  acid-forming  element  it  resembles 
chromium,  forming  salts  of  the  formulas  M2UO4  and  M2U2O?,  but  ordi- 
narily it  plays  the  part  of  a  base-forming  element.  In  this  capacity  it 
forms  a  great  variety  of  salts,  the  best-known  of  which  are  a  series 
in  which  the  radical  UO2,  known  as  uranyl,  acts  as  a  divalent  metal. 
Examples  of  these  are  uranyl  sulfate  (UO2SO4  •  3  H2O),  uranyl  nitrate 
(U02(N08)2-  6  H20),  and  uranyl  acetate  (UO2(C2H8O2)2  •  2  H2O). 

Compounds  of  uranium  are  used  in  glass  making  to  produce  a 
yellowish-green,  fluorescent  glass,  and  in  china  painting  to  secure  a 
velvety  black  color. 

Radioactivity  of  uranium.  In  1896  the  French  physicist  Becquerel 
discovered  that  uranium,  as  well  as  its  compounds,  possesses  a  property 
which  is  called  radioactivity.  A  photographic  plate  wrapped  in  black 
paper  and  placed  near  these  substances  is  affected  as  though  exposed 
to  light.  A  charged  electro-  A^>  A- 

scope  is  rapidly  discharged 
when  in  the  neighborhood  of 
any  of  them,  showing  that  the 
ah*  all  about  them  is  made 
a  conductor. 

Fig.  156    represents   a  simple 
form  of  aluminium-leaf  electron!-        FIG.  156  FIG.  157 

eter,    the    leaves    assuming    the 

position  indicated  at  B  when  an  electric  charge  is  communicated  to  the  knob  A. 
When  a  substance  containing  uranium  (Fig.  157,  C)  is  brought  near  the  knob, 
the  charge  is  rapidly  lost,  and  the  leaves  collapse  as  shown  at  B. 


532  GENERAL  CHEMISTRY 

In  the  case  of  pure  compounds  the  effect  is  proportional  to  the 
quantity  of  uranium  present  in  the  salts,  and  is  independent  of  their 
chemical  character.  The  minerals  containing  uranium,  notably  pitch- 
blende, are  as  much  as  four  times  as  radioactive  as  the  pure  metal. 
This  suggested  that  the  property  of  radioactivity  might  possibly  be 
due  to  some  other  substance  contained  in  the  ore  and  carried  over 
into  uranium  compounds  as  an  impurity. 

Discovery  of  radium.  Accordingly,  Monsieur  and  Madame  Curie 
made  a  careful  study  of  pitchblende,  in  the  hope  of  discovering  some 
unknown  element  in  it.  They  found  that  the  barium  chloride  obtained 
from  the  mineral  (amounting  to  about  30  pounds  per  ton)  was  exceed- 
ingly radioactive,  and  by  a  long-continued  process  of  fractional  crystal- 
lization they  isolated  a  minute  quantity  of  the  chloride  and  bromide  of 
a  new  element,  to  which  they  gave  the  name  radium.  These  compounds 
are  about  three  million  times  as  radioactive  as  uranium.  The  spectrum 
of  this  element,  its  atomic  weight  (226.4),  and  its  chemical  properties 
place  it  in  the  second  group  as  a  homologue  of  barium,  which  «it  very 
closely  resembles.  The  metal  itself,  isolated  by  Madame  Curie  in  1910, 
is  very  similar  to  barium  in  its  reactions  and  appearance. 

Disintegration  of  radium.  The  fact  which  gives  radium  its  peculiar 
interest  is  that,  although  it  is  a  well-characterized  element,  it  is  un- 
stable. It  is  slowly  undergoing  a  change  which  results  in  its  own  dis- 
appearance and  in  the%  emission  of  three  types  of  rays,  designated  as 
alpha  (a),  beta  (/3),  and  gamma  (7)  rays.  The  alpha  rays  are  posi- 
tively charged ;  they  appear  to  consist  of  material  of  the  general  weight 
of  helium  atoms,  and  are  thrown  off  with  high  velocity.  By  many  they 
are  believed  to  be  identical  with  charged  helium  atoms,  since  this  gas 
is  formed  when  alpha  rays  are  given  off.  The  beta  rays  are  negatively 
charged,  move  with  a  velocity  nearly  as  great  as  that  of  light,  and 
appear  to  be  identical  with  electrons  and  cathode  rays.  The  gamma 
rays  seem  to  be  pulsations  in  the  ether  similar  to  Rontgen  rays.  It 
is  estimated  that  in  approximately  fifteen  hundred  years  one  half  of 
any  given  quantity  of  radium  will  have  decomposed  in  this  way. 
Since  radium  is  found  in  ores  which  there  is  every  reason  to  believe 
are  millions  of  years  old,  it  is  clear  that  it  must  be  forming  from  some- 
thing else  and  is  at  the  same  time  undergoing  transition  into  some  other 
permanent  elements. 

Uranium  the  source  of  radium.  The  combined  work  of  numerous 
investigators  has  demonstrated  the  fact  that  uranium  is  the  parent 


THE  VANADIUM  AND  MOLYBDENUM  FAMILIES      533 

substance  from  which  radium  is  derived.  This  element  is  decomposing 
at  a  rate  which  will  accomplish  the  decomposition  of  one  half  of  a  given 
quantity  of  the  element  in  about  six  billion  years.  The  primary  sub- 
stance formed  in  this  decemposition  was  isolated  by  Crookes  and  named 
uranium  X.  This  rapidly  gives  place  to  a  very  transitory  material 
called  ionium,  which  in  turn  forms  radium.  In  all  a  series  of  twelve 
successive  products  has  been  traced,  most  of  which  last  but  a  short 
time,  their  life  being  measured  in  days  or  even  in  minutes.  In  some 
cases  the  transition  from  one  to  the  next  is  accompanied  by  the  escape 
of  helium,  in  others  by  the  expulsion  of  electrons,  in  still  others  by 
both.  The  series  is  as  follows : 

1.  Uranium  5.  Radium  Emanation  (Niton)  9.  Radium  D 

2.  Uranium  X  6.  Radium  A  10.  Radium  E 

3.  Ionium  7.  Radium  B  11.  Radium  F 

4.  Radium  8.  Radium  C  12.  Polonium 

The  end  products.  While  helium,  electrons,  and  radiant  energy  are 
all  given  off  in  these  transitions,  the  major  portion  of  the  original 
uranium  remains  to  be  accounted  for,  and  the  question  arises,  What 
finally  becomes  of  the  uranium  at  the  end  of  the  series  of  transforma- 
tion ?  The  minute  quantity  of  radium  available  for  study,  and  the 
extreme  slowness  with  which  this  product  decomposes,  render  a  direct 
experimental  answer  to  this  question  impossible.  There  are  many  in- 
direct lines  of  evidence  which  have  made  it  seem  probable  that  some 
of  the  familiar  elements  of  lower  atomic  weight,  particularly  lead,  rep- 
resent these  final  products.  It  is  a  significant  fact  that  lead  is  always 
present  in  uranium  ores,  even  when  these  are  found  in  regions  in  which 
no  deposits  of  lead  occur. 

Some  intermediate  substances.  Ramsay  and  Soddy  first  definitely 
proved  that  helium  is  a  permanent  product  of  the  decomposition  of 
radium  emanation ;  it  is  probably  formed  in  a  number  of  other  transi- 
tions. The  radium  emanation  is  a  gas  which  was  first  discovered  by 
Rutherford  and  which  is  now  designated  as  niton.  It  is  condensed  to 
a  solid  by  liquid  air  and  boils  at  about  —  150°.  It  finds  a  place  in  the 
O  group  of  inactive  elements,  its  atomic  weight  being  222.4.  Polo- 
nium was  isolated  by  Madame  Curie  and  named  in  honor  of  her  native 
country,  Poland.  In  many  respects  it  resembles  bismuth  and  tellurium. 

The  supply  of  radium.  It  will  be  seen  that  the  quantity  of  radium  ex- 
isting in  a  mineral  represents  an  equilibrium  between  the  rate  at  which 
it  forms  from  uranium  and  that  at  which  it  in  turn  decomposes.  This 


534  GENERAL  CHEMISTRY 

results  in  its  being  present  in  all  uranium-bearing  minerals  and  in  a 
fixed  ratio  to  the  uranium,  this  ratio  being  1  part  of  radium  to  about 
3,000,000  parts  of  uranium.  There  is  therefore  no  probability  of  the 
discovery  of  any  more  concentrated  ore  of  radium.  In  1  ton  of  pitch- 
blende (60  per  cent  uranium)  there  is  about  0.2  g.  of  the  element. 
Radium  is  at  present  valued  at  about  $100  per  milligram,  or  over 
$2,500,000  per  ounce,  and  there  are  a  number  of  factories  actively 
engaged  in  producing  it. 

Energy  of  radium.  Apparently,  the  rate  at  which  radium  decomposes  cannot 
be  changed  by  any  means  within  the  reach  of  experiment.  Its  decomposition  is 
entirely  spontaneous,  and  the  evolution  of  energy  in  its  progress  is  enormous. 
The  heat  given  off  by  a  gram  of  radium  is  estimated  at  a  little  over  100  cal. 
per  hour,  and  this  is  maintained  throughout  its  entire  life.  Sir  William  Ramsay 
estimates  that  if  the  energy  of  a  ton  of  radium  could  be  evenly  liberated  during 
a  period  of  thirty  years,  it  would  be  sufficient  to  propel  a  ship  of  15,000  tons 
displacement,  driven  by  15,000  horse-power  engines  at  a  rate  of  15  knots  an  hour 
for  the  entire  thirty  years.  Such  a  quantity  of  energy  is  immeasurably  greater 
than  that  which  can  be  obtained  by  chemical  reactions  from  masses  of  matter 
of  like  weight.  This  calculation  gives  us  some  idea  of  the  small  fraction  of  the 
total  energy  of  an  element  which  is  liberated  during  chemical  action,  for  in  its 
ordinary  reactions  radium  does  not  give  off  any  more  heat  than  many  another 
element.  The  steady  evolution  of  heat  by  radioactive  substances  is  held  by  some 
to  be  a  very  plausible  explanation  of  the  origin  of  the  interior  heat  of  the  earth, 
and  possibly  that  of  the  sun  as  well. 

Effects  produced  by  radium.  By  virtue  of  the  energy  liberated,  or  of 
the  particles  shot  off  at  tremendous  velocities,  radium  exerts  a  great 
influence  upon  other  substances.  It  affects  silver  salts  as  exposure  to 
light  does ;  it  produces  fluorescence  in  many  substances,  such  as  barium 
platinocyanide,  zinc  sulfide,  the  diamond,  and  a  number  of  other  gems  ; 
it  discolors  the  glass  of  bottles  in  which  it  is  preserved,  apparently  by 
inducing  chemical  changes  in  the  glass ;  it  sterilizes  seeds  and  kills 
many  microorganisms,  or  at  least  diminishes  their  vitality.  It  is  this 
last  property  which  suggests  its  use  in  medicine  as  a  cure  for  malig- 
nant diseases  such  as  cancer,  and  experiments  are  being  actively 
pushed  to  ascertain  its  value  in  this  connection.  The  water  of  many 
springs  is  radioactive,  on  account  of  materials  held  in  solution,  and  it 
is  thought  that  these  too  may  have  some  medicinal  properties. 

Other  radioactive  elements.  Very  shortly  after  the  discovery  of  the 
radioactivity  of  uranium  compounds  it  was  found  that  those  of  thorium 
are  also  radioactive.  The  rate  of  disintegration  of  thorium  is  much 
slower  than  that  of  uranium,  and  less  is  known  of  the  products 


THE  VANADIUM  AND  , MOLYBDENUM  FAMILIES      535 

formed.  At  least  eight  of  these  have  been  discovered,  and  in  character 
they  resemble  the  corresponding  uranium  products.  Both  potassium 
and  rubidium  have  been  shown  to  be  feebly  radioactive,  and  it  may 
be  that  to  some  degree  the  property  is  possessed  by  all  elements. 

These  facts  at  once  suggest  a  great  many  questions  as  to  the  real 
nature  of  the  bodies  that  we  call  elements.  Some  of  them  are  appar- 
ently in  process  of  evolution,  or  rather  degradation,  from  one  state  to 
another ;  Ramsay  has  suggested  the  outline  of  a  hypothesis  by  which, 
in  time,  the  facts  of  radioactivity  and  many  of  the  relations  of  the 
periodic  law  may  be  brought  together  into  a  general  conception  of 
the  formation  of  the  elements  of  smaller  atomic  weight  from  those 
having  larger  atomic  weights,  with  constant  loss  of  energy. 


CHAPTER  XXXIX 

GOLD  AND  THE  PLATINUM  FAMILY 


ATOMIC 
WEIGHT 

DENSITY 

MELTING 
POINT 

HIGHEST 
OXIDE 

HIGHEST 
CHLORIDE 

Ruthenium  (Ru)  
Rhodium  (Rh) 

101.7 
102  9 

12.3 
1244 

1900.° 
1907° 

RuO4 
RhO 

RuCl4 
RhCl3 

Palladium  (Pd)  
Osmium  (Os)  
Iridium  (Ir)  

106.7 
190.9 
193.1 

11.4 
22.5 
22.41 

1549.° 
2200.° 
2300.° 

Pd02 
OsO4 
IrO2 

PdCl4 
OsCl, 
IrCl4 

Platinum  (Pt)  .  .  . 

195.2 

21.5 

1755  ° 

PtO 

PtCl4 

Gold(Au)  -  / 

197.2 

-  19.32 

1062.4° 

Au203 

AuCl3 

General.  The  periodic  arrangement  places  gold  along  with  copper 
and  silver  in  the  first  group  of  elements,  but  it  is  much  more  closely 
related  to  the  platinum  metals  and  will  be  considered  along  with  them. 

The  eighth  group  in  the  periodic  arrangement  differs  very  much  in 
character  from  the  others,  and  is  made  up  of  three  sets  of  elements, 
each  consisting  of  three  members.  The  first  of  these,  comprising  iron, 
cobalt,  and  nickel,  has  been  described  in  a  previous  chapter.  The  six 
members  of  the  other  two  sets,  including  ruthenium,  rhodium,  palla- 
dium, with  atomic  weights  of  approximately  100,  and  osmium,  iridium, 
and  platinum,  with  atomic  weights  lying  near  the  value  200,  are  very 
closely  related  to  each  other  and  are  known  collectively  as  the  plat- 
inum metals. 

Properties  and  conduct.  In  physical  properties  these  elements  possess 
in  a  high  degree  the  characteristics  of  metals.  For  the  most  part  they 
are  malleable  and  ductile,  are  good  conductors,  crystallize  in  the  regular 
system,  and  have  high  melting  points.  In  chemical  conduct  they  have 
little  affinity  for  other  elements,  and  their  compounds  are  easily  decom- 
posed by  heat,  leaving  a  residue  of  the  pure  metal.  While  each  of 
them  forms  a  number  of  oxides,  these  are  very  unstable,  and  the 
same  property  characterizes  all  of  their  oxygen  compounds.  Their 
hydroxides  are  nearly  all  amphoteric,  but  as  bases  they  form  very 
few  salts  of  oxygen  acids,  while  as  acids  their  salts  are  very  easily 
hydrolyzed. 

536 


GOLD  AND  THE  PLATINUM  FAMILY  537 

Their  more  stable  compounds  are  chiefly  halogen  derivatives,  par- 
ticularly chloro-acids  and  their  salts,  together  with  complex  cyanides 
and  ammonia  compounds.  A  detailed  study  of  these  complexes  is 
beyond  the  scope  of  this  book,  and  most  of  the  individual  chemical 
characteristics  of  the  several  elements  will  be  passed  over,  with  only 
brief  mention  of  a  few  reactions  of  special  importance. 

GOLD 

Occurrence.  From  the  earliest  times  gold  has  been  known  as  a  pre- 
cious metal.  It  was  called  aurum  by  the  Romans,  and  from  this  name 
the  symbol  Au  is  derived.  For  the  most  part  it  is  found  in  nature  in 
the  native  state,  either  embedded  in  quartz  veins  or  as  grains  or  large 
nuggets  in  the  heavy  sands  derived  from  them.  In  this  state  it  is 
usually  alloyed  with  smaller  quantities  of  other  metals,  such  as  silver, 
copper,  and  lead.  In  combination  it  occurs  as  a  constituent  of  a  num- 
ber of  minerals,  nearly  all  of  which  contain  tellurium,  together  with 
silver  and  some  copper.  A  little  gold  telluride  is  apparently  dissolved 
in  the  sulfides  of  many  other  metals,  such  as  those  of  copper,  lead, 
silver,  and  iron,  so  that  gold  is  obtained  as  a  by-product  in  the  refin- 
ing of  these  metals.  The  Transvaal  is  the  largest  producer  of  gold; 
the  United  States  produces  over  one  fifth  of  the  world's  output.  The 
total  output  has  increased  rapidly  during  the  past  years,  and,  since 
little  gold  is  permanently  lost  in  the  arts,  the  total  stock  has  very 
greatly  increased  —  a  fact  which  has  ha<J  a  considerable  economic  impor- 
tance, since  gold  is  the  standard  by  which  other  values  are  measured. 

Extraction.  The  extraction  of  gold  is  accomplished  in  a  number  of 
ways,  depending  upon  the  character  of  the  deposit.  In  placer  mining 
the  gold-bearing  sand  is  washed  by  a  current  of  water  which  is  so 
regulated  that  particles  of  light  weight  are  swept  away,  while  the 
heavier  gold  is  obtained  as  a  deposit.  In  hydraulic  mining  the  earth 
and  sand  are  swept  into  sluices  by  powerful  streams  of  water  operated 
by  pumps.  In  quartz  mining  the  quartz  is  stamped  to  powder  and  is 
then  washed  over  copper  plates,  the  surfaces  of  which  have  been  amal- 
gamated. The  particles  of  gold  stick  to  the  mercury  or  dissolve  in  it, 
the  gold  being  recovered  by  distillation.  In  other  cases,  especially 
when  the  gold  is  in  very  fine  powder  or  in  chemical  combination, 
chemical  reactions  are  employed.  In  the  cyanide  process  (often  used 
to  work  over  the  tailings  from  the  amalgamation  process)  the  gold- 
bearing  material  is  digested  with  a  dilute  solution  of  potassium 


538  GENERAL  CHEMISTRY 

cyanide,  with  free  access  of  air.  The  gold  dissolves  to  form  a  com- 
plex cyanide,  from  which  it  can  be  precipitated  by  metallic  zinc  or 
by  electrolysis : 

4  Au  +  8  KNC  +  02  +  2  H2O  =  4  KAu(NC)2  +  4  KOH 
2  KAu(NC)2  +  Zn  =  K2Zn(NC)4  +  2  Au 

In  the  chlorination  process  the  ore  is  treated  with  chlorine,  which  con- 
verts the  gold  into  the  soluble  trichloride  AuCl3.  It  is  recovered  from 
this  solution  by  suitable  precipitants.  The  treatment  of  lead  and  silver 
ores  containing  gold,  as  well  as  the  separation  of  gold  from  silver,  has 
already  been  described  (p.  495). 

Properties.  Gold  is  a  yellow  metal  of  density  19.32.  It  melts  at 
about  the  same  temperature  as  copper  (1062.4°)  and  boils  at  approxi- 
mately 2500°.  It  is  about  as  soft  as  silver,  is  a  good  conductor  of 
electricity,  and  is  the  most  ductile  and  malleable  of  all  the  metals.  It 
forms  alloys  with  most  of  the  metals,  and  its  uses  are  too  familiar  to 
require  description.  In  a  pure  condition  gold  is  too  soft  to  be  used 
for  jewelry  and  coinage,  and  for  such  purposes  it  is  always  alloyed 
with  copper  or  silver.  The  fineness  of  gold  is  usually  expressed  in 
terms  of  carats,  24-carat  gold  being  pure,  while  18-carat  (75  per  cent) 
is  the  grade  used  for  the  best  jewelry.  For  coinage  a  90  per  cent  alloy 
is  used. 

Chemical  conduct.  Gold  is  not  attacked  by  any  one  of  the  common 
acids.  It  is  easily  dissolved  by  solutions  containing  free  chlorine 
or  bromine,  by  aqua  regia,  and  by  solutions  of  potassium  cyanide  in 
the  presence  of  air.  Fused  alkalies  also  corrode  the  metal,  with  the 
formation  of  aurates  of  the  general  type  KAuO2. 

In  its  compounds  gold  acts  either  as  a  univalent  or  as  a  trivalent 
element.  The  aurous  compounds  recall  the  cuprous  compounds  most 
strongly.  They  oxidize  with  great  ease,  and  sometimes  spontaneously 
decompose,  forming  auric  compounds  and  free  gold : 

3  AuCl  =  2  Au  +  AuCl8 

Auric  hydroxide  (Au  (OH)g)  yields  two  series  of  compounds.  The  one 
is  represented  by  such  salts  as  the  chloride  (AuCl3),  while  the  other 
consists  of  aurates  such  as  potassium  aurate  (KAuO2).  Gold  forms  a 
number  of  sulfides,  such  as  the  monosulfide  (Au2S)  and  the  trisulfide 
(Au2S3).  Like  the  sulfides  of  arsenic,  antimony,  and  tin,  these  are 
soluble  in  ammonium  polysulfide,  with  the  formation  of  thio  salts. 


GOLD  AND  THE  PLATINUM  FAMILY  539 

Complex  compounds.  The  great  majority  of  the  compounds  of  gold 
are  complex  compounds,  such  as  cyanides  and  ammonia  derivatives. 
Among  the  most  important  is  chlorauric  acid  (HAuCl4),  formed  by 
dissolving  gold  in  aqua  reyia.  The  sodium  salt  (NaAuCl4  •  2H2O)  is 
used  as  a  toning  reagent  in  photography.  The  cyanides  KAu(NC)2 
and  KAu  (NC)4  are  used  as  electrolytes  in  gold  plating.  One  of  the 
most  characteristic  reactions  of  gold  compounds  is  the  production  of 
a  purple  precipitate  known  as  the  purple  of  Cassius,  when  a  solution 
containing  stannous  chloride  and  stannic  acid  is  warmed  with  a  very 
dilute  solution  of  a  gold  salt.  Apparently  the  gold  salt  is  reduced  by 
the  stannous  chloride,  and  the  metallic  gold  is  precipitated  as  a  colloidal 
substance,  along  with  colloidal  stannic  acid. 

PLATINUM 

History  and  occurrence.  The  element  platinum  owes  its  name  to 
the  Spanish  word  platina,  which  is  a  diminutive  of  the  word  plata, 
meaning  "silver."  It  was  first  described  with  some  care  by  Brownrigg 
in  1750.  Platinum  occurs  in  nature  alloyed  with  various  other  similar 
metals  (to  be  described  shortly),  as  well  as  with  gold,  copper,  and  iron. 
It  is  found  chiefly  in  the  Ural  Mountains  in  Russia,  and  in  Brazil, 
Mexico,  Colombia,  and  the  Pacific  coast  states,  occurring,  like  gold,  in 
heavy  sands,  associated  with  magnetite,  rutile,  diamond,  and  quartz, 
and  constituting  from  50  to  80  per  cent  of  the  crude  alloy.  A  single 
well-defined  mineral  compound  is  known;  namely,  sperrylite  (PtAs2). 
Some  platinum  is  now  recovered  from  the  electrolytic  mud  of  the  gold 
and  silver  refineries.  The  separation  of  the  platinum  from  the  other 
metals  with  which  it  is  alloyed  is  a  very  complicated  process  and  cannot 
be  described  here; 

Properties.  Platinum  is  a  grayish- white  metal  of  density  21.5,  which 
melts  at  1755°.  It  is  very  malleable  and  ductile  and  not  very  hard 
when  pure.  Under  ordinary  conditions  it  is  not  attacked  by  the  com- 
mon acids,  though  hot,  concentrated  sulfuric  acid  slowly  dissolves  it. 
It  is  quite  permanent  in  the  air,  and  neither  oxygen  nor  water  vapor 
attacks  it  even  at  red  heat.  Free  chlorine  dissolves  it,  forming  plati- 
num tetrachloride  (PtCl4).  Aqua  regia  converts  it  into  chloroplatinic 
acid  (H2PtCl6).  Fused  alkalies  corrode  it,  forming  platinates,  espe- 
cially in  the  presence  of  oxidizing  agents  such  as  nitrates.  It  alloys 
readily  with  many  easily  reducible  metals  and  is  therefore  attacked 
or  dissolved  when  heated  with  compounds  of  such  metals  together 


540  GENERAL  CHEMISTRY 

with  a  reducing  agent.  Since  hydrogen  gas  readily  passes  through 
sheet  platinum  when  the  latter  is  hot,  it  is  not  safe  to  heat  any  easily 
reducible  oxide  in  a  platinum  crucible  in  a  gas  flame. 

Uses.  Platinum  is  well  adapted  to  many  uses  and  would  find  many 
applications  if  it  were  not  so  costly.  Its  resistance  to  chemical  action 
and  its  high  melting  point  make  it  invaluable  for  laboratory  utensils. 
Its  low  position  in  the  electromotive  series  makes  it  well  adapted 
for  use  as  electrodes.  It  is  a  moderately  good  conductor  of  electricity, 
and  as  its  coefficient  of  expansion  with  heat  is  about  the  same  as  that 
of  soft  glass,  it  is  particularly  well  fitted  for  sealing  through  glass 
tubes  as  an  electric  conductor  —  for  example,  in  electric-light  bulbs. 
Its  catalytic  effect  on  many  chemical  reactions  suggests  its  use  in  a 
number  of  chemical  industries,  as  in  making  sulfuric  acid.  Lastly,  it 
has  become  the  fashion  in  jewelry,  especially  for  the  setting  of  bril- 
liants and  for  watch  chains,  and  at  present  these  uses  absorb  a  large 
part  of  the  production. 

Compounds  of  platinum.  Like  tin,  platinum  acts  either  as  a  divalent 
or  as  a  tetravalent  element,  and  the  hydroxides  corresponding  to  each 
valence  can  act  either  as  acids  or  as  bases.  Some  trivalent  compounds 
are  also  known.  In  general,  the  oxygen  derivatives  of  platinum,  in- 
cluding the  oxygen  acids,  as  well  as  the  platinates  and  platinites,  are 
unstable  and  easily  decomposed  by  heat.  Its  best-known  compounds 
are  its  halogen  derivatives,  especially  the  salts  of  chloroplatinic  acid 
(H2PtCl6).  This  free  acid,  which  forms  reddish,  deliquescent  crystals 
of  the  composition  H2PtCl6  •  6H2O,  is  obtained  by  dissolving  platinum 
in  aqua  regia.  Its  potassium  and  ammonium  salts  are  nearly  insolu- 
ble in  dilute  alcohol  and  are  easily  separated  from  the  sodium  salt, 
which  is  freely  soluble.  Barium  platinocyanide  (BaPt(NC)4  •  4  H2O) 
phosphoresces  under  the  influence  of  Rontgen  rays  and  is  of  impor- 
tance in  X-ray  photography.  There  are  a  great  many  ammonia 
derivatives,  which  exhibit  many  interesting  and  complicated  relations 
among  themselves. 

Other  platinum  metals.  The  other  platinum  metals  —  palladium, 
rhodium,  ruthenium,  osmium,  and  iridium  —  are  all  gray  metals  and 
have  very  high  melting  points  and  densities.  They  are  found  alloyed 
with  native  platinum.  In  chemical  conduct  they  are  like  gold  and 
platinum.  Each  one  forms  several  series  of  compounds,  in  which  the 
valence  is  either  2,  3,  or  4 ;  each  gives  a  variety  of  oxygen  com- 
pounds of  rather  unstable  character,  and  each  forms  more  stable 


GOLD  AND  THE  PLATINUM  FAMILY  541 

chloroacids  and  complex  derivatives  of  ammonia  and  cyanogen.  Nearly 
all  of  these  compounds  are  highly  colored  in  a  great  variety  of  tints. 

Palladium.  Palladium  is  the  most  abundant  one  of  the  five.  It  is 
obtained  from  native  platinum  and  also  from  the  electrolytic  muds  in 
the  refining  of  certain  nickel  ores.  It  is  a  soft  metal  closely  resembling 
platinum  in  appearance,  but  of  less  density  and  of  lower  melting  point. 
It  is  more  pronouncedly  basic  than  the  others,  dissolving  in  concen- 
trated acids,  especially  in  nitric  acid,  and  forming  a  number  of  divalent 
and  moderately  stable  oxygen  salts,  such  as  the  nitrate  (Pd(NO3)2)  and 
the  sulfate  (PdSO4).  As  an  elementary  substance  its  most  remark- 
able property  is  its  capacity  for  absorbing  large  volumes  of  hydrogen. 
When  the  metal  is  finely  divided,  this  may  amount  to  as  much  as  eight 
hundred  volumes,  and  an  even  larger  volume  is  absorbed  when  an  elec- 
trode is  covered  with  the  spongy  metal  and  is  made  the  cathode  in  the 
electrolysis  of  a  dilute  acid  solution.  The  hydrogen  so  absorbed  is  in 
a  very  active  state,  showing  the  reactions  appropriate  to  its  position  in 
the  electromotive  series.  For  example,  such  a  charged  electrode,  dipped 
into  a  solution  of  copper  salt,  at  once  precipitates  copper. 

Palladium  is  used  in. making  graduated  vernier  scales  and  as  a  sub- 
stitute for  platinum  in  jewelry.  While  it  commands  the  same  price 
as  platinum,  it  is  only  half  as  heavy,  so  that  a  given  weight  will  go 
twice  as  far.  It  is  also  used  as  a  catalyzer  in  industrial  reactions, 
as  in  the  reduction  of  certain  fats  and  oils. 

Rhodium.  Rhodium  resembles  aluminium  in  appearance.  Of  all 
these  metals,  it  is  the  most  easily  attacked  by  free  chlorine,  but  it  is 
exceedingly  resistant  to  the  action  of  acids. 

Ruthenium.  Ruthenium  is  hard  and  brittle  and  is  dark  gray  or 
nearly  black  in  color.  It  forms  a  variety  of  oxygen  compounds,  the 
formulas  of  which  recall  those  of  manganese  —  for  example,  RuO2, 
K2RuO4,  and  KRuO4.  It  also  forms  a  volatile  oxide  of  the  formula 
RuO4.  This  is  produced  by  the  oxidizing  action  of  aqua  regia,  and  is 
possibly  formed  by  heating  the  metal  in  the  air.  It  is  a  yellow  com- 
pound which  melts  at  about  26°,  boils  at  about  100°,  and  is  volatile 
with  steam. 

Osmium.  Osmium  is  interesting  as  being  the  heaviest  of  all  known 
substances,  having  a  density  of  22.5.  It  is  very  hard  and  infusible,  and 
is  chiefly  acid-forming  in  character.  It  forms  an  oxide  (OsO4)  which 
resembles  the  corresponding  oxide  of  ruthenium.  This  oxide  is  volatile 
with  steam  and  is  formed  by  the  action  of  aqua  regia  upon  osmium. 


542  GENERAL  CHEMISTRY 

It  is  often  present  in  the  steam  arising  when  solutions  of  crude  platinum 
salts  are  boiled,  and  has  a  very  irritating  and  unpleasant  odor.  It  is 
called  osmic  acid,  though  it  has  no  acid  properties,  nor  does  it  form 
an  acid  with  water.  The  name  was  given  on  account  of  its  corrosive 
action,  which  it  owes  to  oxidizing  and  not  to  acid  properties.  In  acting 
as  an  oxidizing  agent  it  deposits  metallic  osmium  in  very  finely  divided 
form,  and  this  is  a  great  irritant  when  deposited  in  sensitive  tissues. 
It  is  used  in  biological  laboratories  as  a  hardening  agent  and  as  a  stain 
in  the  preparation  of  microscopic  sections. 

Iridium.  This  element  owes  its  name  to  the  variety  of  colors  seen 
during  its  chemical  transformations.  It  is  a  silver-white  metal,  hard 
and  brittle.  It  is  often  present  in  the  platinum  of  laboratory  vessels, 
making  the  latter  harder  and  less  subject  to  chemical  corrosion.  The 
standard  meter  bar  preserved  at  Sevres  is  an  alloy  of  platinum  and 
iridium.  The  residue  obtained  when  native  platinum  is  digested  with 
aqua  regia  is  essentially  an  alloy  of  iridium  and  osmium.  It  can  be 
worked  up  into  a  very  hard  alloy  called  iridosmine,  which  is  used  as 
a  material  for  pointing  gold  pens. 


INDEX 


Absolute  scale  of  temperature,  48 

Accumulator,  610 

Acetates,  305 

Acetic  acid,  305 ;  glacial,  305 

Acetylene,  295,  423 ;  structure  of,  293 

Acid-forming  elements,  234,  382 

Acids,  137,  148;  action  of,  on  elements, 
379 ;  basicity  of,  149  ;  complex,  529 ; 
condensed,  346;  formed  by  water  and 
oxides,  62 ;  nomenclature  of,  159 ;  nor- 
mal, 377;  poly-,  346;  preparation  of, 
from  salts,  200 ;  relative  ionization  of, 
155 ;  strength  of,  149 

Active  mass,  194 

Adsorption,  281 

Affinity,  191 

Affinity  constant,  193 

Afterdamp,  331 

Agent,  catalytic,  28 ;  oxidizing,  23,  471 ; 
reducing,  43 

Air,  114 ;  a  mixture,  119 ;  carbon  dioxide 
in,  118 ;  constancy  of  composition  of, 
119;  constituents  of,  114;  dust  parti- 
cles in,  115;  impure,  120;  liquefaction 
of,  78;  microorganisms  in,  116;  deter- 
mination of  nitrogen  in,  117;  determi- 
nation of  oxygen  in,  116 ;  properties  of, 
121;  solubility  of,  119;  variation  in 
composition  of,  118 ;  volumetric  analysis 
of,  116 ;  water  vapor  in,  118 

Air  saltpeter,  182 

Alabaster,  421 

Albite,  347,  447 

Alchemists,  29 

Alcohol,  absolute,  302;  denatured,  303; 
ethyl,  301 ;  methyl  (wood),  301 

Alcoholic  liquors,  303 

Alcohols,  301 

Aldehydes,  304 

Alkali  metals,  391 

Alkalies,  150 

Alkaline  earth  family,  413 

Alloys,  374,  375,  480 

Alum,  445,  449  ;  ammonium,  446 ;  chrome, 
522;  iron,  470;  potassium,  446 

Aluminic  acid,  443 

Aluminium,  439 ;  a  reducing  agent,  384  ; 
compounds  of,  442  ;  preparation  of,  440 ; 
properties  of,  441  ;  uses  of,  441 

Aluminium  bronze,  442 

Aluminium  carbide,  447 

Aluminium  carbonate,  446 


Aluminium  chloride,  444 

Aluminium  group,  439 

Aluminium  hydroxide,  443 

Aluminium  oxide,  443 

Aluminium  silicates,  447 

Aluminium  sulfate,  445 

Aluminium  sulfide,  445 

Alundum,  433 

Alunite,  446 

Amalgams,  374,  487 

Amblygonite,  392 

Amethyst,  344 ;  oriental,  443 

Amido  group,  172 

Amino  group,  172 

Ammonia,  168;  action  of,  on  chlorine,  251; 
action  of,  on  Salts,  172,  492  ;  chemical 
conduct  of,  170;  commercial  produc- 
tion of,  170,  424;  composition  of,  173; 
equilibrium  of,  with  hydrogen  and 
nitrogen,  195;  properties  of,  170;  solu- 
bility of,  in  water,  123;  uses  of,  174 

Ammonia  water,  169 

Ammoniacal  liquor,  170,  323 

Ammoniates,  492 

Ammonium,  172,  410 ;  compounds  of, 
410 ;  equilibrium  of,  in  solution,  125 

Ammonium  amalgam,  410 

Ammonium  bicarbonate,  412 

Ammonium  bisulfate,  412 

Ammonium  bromide,  411 

Ammonium  carbamate,  412 

Ammonium  carbonate,  412 

Ammonium  chloride,  411 

Ammonium  cyanate,  291 

Ammonium  fluoride,  411 

Ammonium  hydroxide,  172 

Ammonium  iodide,  411 

Ammonium  metavanadate,  527 

Ammonium  molybdate,  529 

Ammonium  nitrate,  412  ;  in  air,  115 

Ammonium  persulfate,  412 

Ammonium  phosphomolybdate,  529 

Ammonium  polysulfides,  412 

Ammonium  sulfate,  412 

Ammonium  sulfides,  411 

Ammonium  sulfostannate,  502 

Ammonolysis,  493 

Amorphous  bodies,  79 ;  solutions  of,  131 

Amphoteric  hydroxides,  372,  378 

Anglesite,  504 

Anhydrides,  188 

Anhydrite,  421 


543 


544 


GENERAL  CHEMISTRY 


Aniline,  297 

Anion,  140 

Anode,  140 

Anthracene,  297 

Antimony,  370 ;  acids  of,  373 ;  alloys  of, 
375 ;  black,  371  ;  compounds  of,  370 ; 
gray,  371 ;  yellow,  371 

Antimony  hydride,  371 

Antimony  hydroxide,  372 

Antimony  pentasulfide,  373 

Antimony  pentoxide,  373 

Antimony  tetroxide,  372 

Antimony  thio  acids,  373 

Antimony  trioxide,  372 

Antimony  trisulfide,  373 

Antimonyl  compounds,  372,  373 

Apatite,  414 

Aqua  ammonia,  169 

Aqua  regia,  257 

Aragonite,  420 

Argon,  108,  110;  spectrum  of,  112 

Arsenic,  364;  antidote  for,  368;  forms 
of,  365 ;  halogen  compounds  of,  366 ; 
Marsh's  test  for,  366 ;  thio  salts  of,  370  ; 
white,  367 

Arsenic  acids,  368,  369 

Arsenic  chloride,  366      - 

Arsenic  disulfide,  369 

Arsenic  hydride,  365 

Arsenic  pentasulfide,  369 

Arsenic  pentoxide,  368 

Arsenic  sulfide,  369 

Arsenic  trioxide,  367,  453 

Arsenic  trisulfide,  369 

Arsenides,  365 

Arsenolite,  364 

Arsenopyrite,  364 

Arsine,  365 

Asbestos,  430 

Atmosphere,  114 

Atomic  theory,  93 ;  value  of,  97 

Atomic  volumes,  curve  of,  239 

Atomic  weights,  90,  316;  accurate  de- 
termination of,  317;  and  molecular 
weights,  96 ;  deduction  of,  from  molec- 
ular weights,  316 ;  relation  of,  to  prop- 
erties, 235 ;  table  of,  12  and  Appendix 
A ;  uses  of,  90 

Atoms,  93 ;  changeability  of,  95,  535 ;  size 
of,  94 

Auer  metal,  452 

Aurates,  538 

Auric  compounds,  538 

Aurous  compounds,  538 

Avogadro's  hypothesis,  310 

Azurite,  484 

Babbitt  metal,  375 

Baking  powders,  alurn,  447  ;   cream  of 

tartar,  402 
Barite,  428 
Barium,  427 
Barium  carbonate,  428 


Barium  chloride,  428 

Barium  chromate,  428 

Barium  hydroxide,  428 

Barium  nitrate,  428 

Barium  oxide,  427 

Barium  peroxide,  70,  427 

Barium  platinocyanide,  540 

Barium  sulfate,  428 

Barium  sulfide,  428 

Base-forming  element,  234,  382 

Bases,  137, 150 ;  acidity  of,  151 ;  action  of, 
on  metals,  38,  379 ;  action  of,  on  non- 
metals,  380 ;  diacid,  151  ;  formed  by 
water  with  oxide,  62  ;  monacid,  151 ; 
nomenclature  of,  159  ;  relative  ioniza- 
tion  of,  155  ;  strength  of,  157  ;  tetracid, 
151  ;  triacid,  151 

Bauxite,  439 

Beer,  303 

Benzene,  297  ;  structure  of,  293 

Benzine,  294 

Berkeland  and  Eyde  process,  177 

Bessemer  converter,  462 

Bessemer  process,  462 

Betts  process,  505 

Bismuth,  374 ;  compounds  of,  375  ;  hy- 
drolysis of  salts  of,  376  ;  oxides  of,  375 

Bismuth  chloride,  376 

Bismuth  nitrate,  376 

Bismuth  trioxide,  376 

Bismuth  subcarbonate,  376 

Bismuth  subnitrate,  376 

Bismuth  sulfide,  376 

Bismuthyl  compounds,  376 

Bisque,  457 

Blast  furnace,  460 

Bleaching,  by  bromine,  261  ;  by  chlorine, 
253 ;  by  hydrogen  peroxide,  72 ;  by 
ozone,  33  ;  by  sulfurous  acid,  214 

Bleaching  powder,  270 

Blister  copper,  479 

Blue  printing,  473 

Bluestone,  483 

Boiler  scale,  433 

Boiling  point,  75  ;  elevation  of,  by  solute, 
134 

Bonds,  163 

Bone  black,  279 

Bone  oil,  280 

Boracic  acid,  352 

Boracite,  351 

Borax,  352 

Bordeaux  mixture,  483 

Boric  acid,  352 

Boric  oxide,  352 

Bornite,  478 

Boron,  351 ;  acids  of,  352  ;  compounds  of, 
352  ;  salts  of  acids  of,  352 

Boyle's  law,  47 

Brass,  480 

Braunite,  516 

Brimstone,  204 

Brin  process,  21,  427 


INDEX 


545 


Britannia  metal,  375 

Bromates,  273 

Bromides,  263 

Bromine,  259  ;  properties  and  chemical 
conduct  of,  261  ;  oxygen  acids  of,  273 

Bronze,  480 ;  aluminium,  480 ;  manga- 
nese, 515 

Butter  fat,  307 

Butyric  acid,  305 

By-products,  255 

Cadmium,  437  ;  compounds  of,  438 

Cadmium  bromide,  438 

Cadmium  chloride,  438 

Cadmium  hydroxide,  438 

Cadmium  iodide,  438 

Cadmium  nitrate,  438 

Cadmium  oxide,  438 

Cadmium  sulfate,  438 

Cadmium  sulfide,  438 

Caesium,  409 

Calamine,  434 

Calcite,  420 

Calcium,  414  ;  compounds  of,  415 

Calcium  bicarbonate,  420 

Calcium  carbide,  423 

Calcium  carbonate,  419 

Calcium  chloride,  418 

Calcium  cyanamide,  424 

Calcium  fluoride,  418 

Calcium  hydroxide,  130,  417 

Calcium  oxalate,  425 

Calcium  oxide,  416 

Calcium  phosphates,  425 

Calcium  silicates,  425 

Calcium  sulfate,  421 

Calcium  sulfide,  418 

Calcium  sulfites,  422 

Calculations,  101  ;  of  volume,  102,  320 

Caliche,  403  ;  extraction  of  iodine  from, 
264 

Calomel,  489 

Calorie,  6 

Calorimeter,  7  ;  bomb,  334 

Caramel,  298 

Carat,  276 

Carbamic  acid,  288 

Carbides,  281 

Carbohydrates,  298 

Carbon,  275 ;  amorphous  forms  of,  277 ; 
chemical  conduct  of,  281  ;  compounds 
of,  281 ;  cycle  of,  in  nature,  121  ;  occur- 
rence of,  275 ;  properties  of,  280 ;  uses 
of,  281 

Carbon  dioxide,  282 ;  absorption  of,  by 
plants,  119,  121^"  critical  constants  of, 
76 ;  heat  of  formation  of,  336 ;  in  air, 
118  ;  preparation  of,  282  ;  reduction  of, 
286  ;  solubility  of,  in  water,  123  ;  uses 
of,  284 

Carbon  disulfide,  289 

Carbon  monoxide,  286 ;  heat  of  forma- 
tion of,  336  ;  reducing  power  of,  287 


Carbon  tetrachloride,  298 

Carbonates,  285  ;  acid,  286  ;  normal,  285; 
solubility  of,  390 

Carbonic  acid,  285 ;  simple  derivatives  of, 
288 

Carbonic  anhydride,  282 

Carbonyl  chloride,  289 

Carbonyls,  476 

Carborundum,  348 

Carnallite,  406 

Casein,  299 

Cassiterite,  500 

Castner's  process,  394 

Catalysis,  28,  192 

Catalytic  agent,  21 

Cathode,  140 

Cation,  140 

Caustic  potash,  405 

Caustic  soda,  396 

Celestite,  426 

Celluloid,  300 

Cellulose,  300 

Cement,  455 

Cementite,  460 

Ceramic  industries,  453 

Cerium,  451 

Cerussite,  504 

Chalcocite,  481 

Chalcopyrite,  478 

Chamber  acid,  219 

Chamber  crystals,  219 

Chamberlain-Pasteur  filter,  58 

Charcoal,  278  ;    absorption  of  gases  by, 
280 ;   animal,  279 

Cheese,  299 

Chemical  action,  10 

Chemical  calculations,  101 

Chemical  combination,  laws  of,  83 

Chemical  conduct,  10 

Chemical  energy,  9,  333 
"Chemical  reactions,  8 

Chemistry,  object  of,  1 

Chile  saltpeter,  403 

Chinaware,  456 

Chlorates,  271 

Chlorauric  acid,  539 

Chloric  acid,  271 

Chloride  of  lime,  270 

Chlorides,  258  ;  solubility  of,  390 

Chlorine,  248 ;  action  of,  on  bases,  380 ; 
\chemical    conduct   of,    251,   252,   253; 
preparation  of,  248,  249,  250 ;  proper- 
ties of,  250  ;  uses  of,  253 

Chlorine  family,  243 

Chlorine  hydrate,  252 

Chlorine  oxides,  268 

Chlorine  water,  252 

Chlorites,  271 

Chloroform,  298 

Chlorophyll,  459 

Chloroplatinic  acid,  539,  540 

Chlorostannic  acid,  503 

Chlorosulfonic  acid,  229 


546 


GENERAL   CHEMISTRY 


Chlorous  acid,  271 

Choke  damp,  330 

Chromates,  523  ;  oxidizing  action  of,  524 

Chrome  alum,  522 

Chrome  iron  ore,  522 

Chrome  red,  510 

Chrome  yellow,  510 

Chromic  acid,  523 

Chromic  anhydride,  523 

Chromic  chlorides,  521 

Chromic  compounds,  521 

Chromic  hydroxide,  521 

Chromic  oxide,  521 

Chromic  sulfates,  522 

Chromite,  519 

Chromites,  522 

Chromium,  519 

Chromium  acids  and  their  salts,  523,  524 

Chromous  compounds,  520 

Chromous  hydroxide,  521 

Chromous  oxide,  520 

Cinnabar,  491 

Citric  acid,  306 

Clay,  448 

Clay  products,  456 

Coal,  278 

Coal  gas,  323 

Coal  oil,  294 

Coal  tar,  294,  323 

Cobalt,  473  ;  compounds  of,  474 

Cobalt  carbonyl,  476 

Cobalt  chloride,  474 

Cobalt  nitrate,  474 

Cobalt  sulfide,  474 

Cobaltammines,  475 

Cobaltite,  474 

Cobaltous  oxide,  474 

Cochineal,  150 

Coins,  composition  of,  480 

Coke,  278 

Colemanite,  351 

Colloidal  suspension,  132 

Colloids,  132 

Columbite,  527 

Columbium,  527  ;  compounds  of,  528 

Combining  weights,  determination  of,  88  ; 

law  of,  86  ;  of  compounds,  88  ;  of  ele- 
ments, 87 
Combustion,  25 ;  discovery  of  nature  of, 

30  ;  effect  of,  on  composition  of  air,  118  ; 

phlogiston  theory  of,  29;  spontaneous,  26 
Compounds,  7  ;  combining  weight  of,  88  ; 

definition  of,   9 ;    general   methods  of 

preparation  of,  386 
Concentration  and  speed  of  reaction,  28, 

192 

Condensation,  heat  of,  76 
Congo  red,  150 

Constant-boiling  solutions,  126,  127 
Contact  process  for  sulfuric  acid,  217 
Copper,   478 ;    ammonia  compounds  of, 

484  ;  blister,  479  ;  metallurgy  of,  478  • 

uses  of,  480 


Copper  arsenide,  365 
Copper  arsenite,  368 
Copper  cyanide,  485 
Copper  ferrocyanide,  134,  473 
Copperas,  467 
Corn  sirup,  300 
Corrosive  sublimate,  80,  491 
Corundum,  443 
Cream  of  tartar,  306 
Cristobalite,  344 
Critical  point,  76 
Critical  pressure,  76 
Critical  temperature,  76 
Crocoite,  519 
Cryolite,  244,  439 
Crystalline  forms,  275 
Crystallization,  water  of,  63 
Crystallography,  systems  of,  81 
Crystals,  81  ;  melting  point  of,  79  ;  struc- 
ture of,  82 
Cupric  acetate,  484 
Cupric  bromide,  484 
Cupric  carbonate,  484 
Cupric  chloride,  484 
Cupric  compounds,  482 
Cupric  hydroxide,  483 
Cupric  nitrate,  484 
Cupric  oxide,  482 
Cupric  sulfate,  130,  483 
Cupric  sulfide,  483 
Cuprite,  481 
Cuprous  acetylide,  482 
Cuprous  bromide,  482 
Cuprous  chloride,  481 
Cuprous  compounds,  481 
Cuprous  cyanide,  482 
Cuprous  iodide,  482 
Cuprous  oxide,  481 
Cuprous  sulfide,  481 
Cyanamide,  424 
Cyanates,  291 
Cyanic,  acid,  291 
Cyanogen,  291 

Daniell  cell,  485 

Deacon's  process,  249 

Decay,  118 

Decomposition,  100 

Definite  composition,  law  of,  84 

Dehydration,  377,  378 

Density,  3 

Developer,  498 

Dewar  flask,  78 

Dextrin,  299 

Dextrose,  299 

Dialysis,  348 

Dialyzer,  348 

Diamond,  275 

Diastase,  299 

Dichromates,  523 

Dichromic  acid,  523 

Diffusion  of  gases,  50 

Dimorphous  substances,  82 


INDEX 


54T 


Dioxides,  507 

Disilicic  acid,  347 

Dissociation,  by  heat,  197  ;  of  electrolytes, 
137 

Distillation,  57  ;  destructive,  280 ;  frac- 
tional, 127  ;  under  diminished  pressure, 
272 

Dithionic  acid,  227 

Dolomite,  432 

Dor6  bars,  495,  505 

Double  bonds,  293 

Double  decomposition,  100,  388,  389 

Dyeing,  444 

Dynamite,  303 

Dysprosium,  451 

Earth's  crust,  composition  of,  14 

Edison  cell,  510 

Efflorescence,  63,  198 

Electric  cells,  485 

Electric  furnace,  385 

Electrode,  140 

Electrode  potential,  158 

Electrolysis,  140 ;  of  sodium  chloride, 
143  ;  of  sodium  sulf  ate,  145 ;  of  sulf  uric 
acid,  144 

Electrolytes,  132  ;  abnormal  solution  con- 
stants of,  139  ;  chemical  action  of,  139  ; 
definition  of,  140  ;  dissociation  of,  138  ; 
relative  ionization  of,  155 

Electromotive  series,  158 

Electrons,  143,  532 

Electroplating,  499 

Electrotyping,  480 

Elements,  11  ;  classification  of,  233  ;  com- 
bining weight  of ,  87  ;  disintegration  of, 
535;  distribution  of,  13;  essential  to  life, 
13  ;  molecular  weight  of,  318  ;  number 
of,  11;  table  of,  12  and  Appendix  A 

Emery,  443 

Emulsion,  128 

Endothermic  compounds,  340 

Endothermic  reactions,  25 

Energy,  4 ;  bound,  332  ;  changes  in,  92  ; 
chemical,  5  ;  conservation  of,  4 ;  free, 
332  ;  liberated  in  stages,  332  ;  measure- 
ment of,  6  ;  total,  332  ;  transformation 
of,  4  ;  varieties  of,  4 

Epsom  salts,  432 

Equations,  92  ;  combining  of,  180 

Equilibrium,  191,  195,  196  ;  and  precipita- 
tion, 200  ;  demonstration  of,  195  ;  fac- 
tors affecting,  196  ;  in  solution,  125 ;  of 
ions,  141,  198 

^Equivalent,  determination  of,  89 
Equivalent  weight,  88 
Erbium,  451 
Esters,  306 
Etching  of  glass,  247 
Ethane,  295 

Ethylene,  295  ;  structure  of,  293 
Eudiometer,  65,  66 
Europium,  451 


Evaporation,  73 

Exothermic  compounds,  340 ;  reaction  of, 

25 

Explosions,  329  ;  dust,  331  ;  mine,  330 
Explosive  mixtures,  329 
Explosives,  331 

Families,  236 

Family  resemblances,  238 

Fatty  acids,  304 

Fehling's  solution,  483 

Feldspar,  347 

Fermentation,  116;  acetic,  305;  alcoholic, 
301  ;  lactic,  306 

Ferric  acid,  472 

Ferric  ammonium  alum,  470 

Ferric  chloride,  469 

Ferric  compounds,  469 

Ferric  f  errocyanide,  473 

Ferric  hydroxide,  469 

Ferric  nitrate,  470 

Ferric  oxide,  469 

Ferric  phosphate,  470 

Ferric  sulfate,  470 

Ferric  sulfocyanate,  470 

Ferricyanic  acid,  473 

Ferrite,  462 

Ferrochromium,  520 

Ferrocyanic  acid,  472 

Ferromanganese,  515 

Ferrosilicon,  342 

Ferrous  ammonium  sulfate,  468 

Ferrous  carbonate,  469 

Ferrous  chloride,  467 

Ferrous  compounds,  466 

Ferrous  ferricyanide,  473 

Ferrous  hydroxide,  467 

Ferrous  oxide,  467 

Ferrous  sulfate,  467 

Ferrous  sulfide,  468 

Ferrovanadium,  526 

Fertilizers,  363 ;  nitrogen,  424;  phosphate, 
363  ;  potassium,  406,  407 

Fire  damp,  295,  330 

Fire  extinguishers,  284 

Fischer's  salt,  475 

Flames,  322  ;  analysis  of,  327  ;  complex, 
326 ;  luminosity  of,  327 ;  oxidizing, 
329  ;  reducing,  329  ;  simple,  326  ;  struc- 
ture of,  326  ;  temperature  of,  328 

Flint,  344 

Fluorapatite,  244 

Fluorides,  247 

Fluorine,  244  ;  properties  of,  246 

Fluorspar,  244,  418 

Fluosilicates,  343 

Fluosilicic  acid,  343,  449 

Fluotitanic  acid,  350 

Fluozirconic  acid,  351 

Fluxes,  353,  362,  384 

Foods,  308,  336 ;  calorific  values  of,  337 ; 
composition  of,  308 

Fool's  gold,  468 


548 


GENERAL  CHEMISTRY 


Formaldehyde,  304 

Formalin,  304 

Formic  acid,  305 

Formula  weights,  91 

Formulas,  90;  calculation  of,  from  mo- 
lecular weights,  315;  calculation  of, 
from  percentages,  91;  molecular,  97; 
structural,  164 

Franklinite,  434 

Fraunhofer  lines,  113 

Freezing  point,  79;  relation  to  concen- 
tration, 133,  310 

Freezing-point  method  of  determining 
molecular  weights,  315 

Fruit  sugar,  300 

Fuel  gases,  322 

Fuels,  336 ;  calorific  values  of,  336 

Fusible  white  precipitate,  493 

Fusion,  heat  of,  79 

Gadolinium,  451 

Galenite,  504 

Gallium,  450 

Garnierite,  475 

Gas,  definition  of,  73;  displacement  of, 
from  solution,  126 ;  natural,  324 ;  oil, 
324 ;  producer,  324 ;  water,  323 

Gas  laws,  46 

Gas  mantles,  324,  512 

Gaseous  state,  45 

Gases,  comparative  composition  of,  325 ; 
compressibility  of,  46 ;  densities  of,  43 ; 
expansibility  of,  45;  liquefaction  of,  77 ; 
properties  of,  45 ;  relation  between 
liquids  and,  73  ;  solubility  of,  in  boiling 
liquids,  126 ;  solution  of,  in  gases,  122 

Gasoline,  294 

Gems,  artificial  preparation  of,  443 

Germanium,  512 

Glass,  453  ;  color  of,  454 ;  ingredients  of, 
453 ;  molding  of,  454 ;  varieties  of,  454 

Glasses,  347 

Glauber's  salt,  399 

Glucinum,  438 

Glucose,  300 

Glycerin,  303 

Gold,  537  ;  compounds  of,  538 

Goldschmidt  reduction  method,  385 

Goldschmidt  welding  process,  442 

Gram,  2 

Gram-molecular  volume,  312 

Gram-molecular  weights,  97 

Grape  sugar,  299 

Graphite,  276 

Group  resemblances,  238 

Gun  metal,  480 

Guncotton,  300 

Gunpowder,  331 ;  smokeless,  300 

Gypsum,  421 

Haemoglobin,  459 
Halogens,  243 
Hard  water,  422 


Hausmannite,  516 

Health,  affected  by  water,  55 

Heat,  and  temperature,  24 ;  in  reversible 
reactions,  338;  law  of  summation  of, 
335 ;  measurement  of,  6 ;  of  changes  in 
state,  333  ;  of  condensation,  76 ;  of  endo- 
thermic  compounds,  334 ;  of  formation, 
337 ;  of  formation  and  decomposition, 
335;  of  fusion,  79;  of  neutralization, 
157;  of  reaction,  334 ;  of  solidification, 
79 ;  of  solution,  334,  338 ;  of  vaporiza- 
tion, 76 

Helium,  109 ;  from  radium,  532 ;  spec- 
trum of,  113 

Hematite,  459,  469 

Hexane,  294 

Homologous  series,  292 

Human  body,  average  composition  of,  13 

Humidity,  relative,  74 

Hyacinth,  351 

Hydrates,  62 

Hydrazine,  174 

Hydriodic  acid,  267 

Hydrobromic  acid,  263 

Hydrocarbons,  292;  properties  of,  295; 
sources  of,  294 ;  substitution  products 
of,  297  . 

Hydrochloric  acid,  257 

Hydrocyanic  acid,  291 

Hydrofluoric  acid,  247 

Hydrogel,  348 

Hydrogen,  35 ;  absorption  of,  by  palla- 
dium, 39,  541 ;  action  of,  on  elements, 
40,  99 ;  chemical  conduct  of,  40 ;  critical 
constants  of,  76  ;  equations  of  prepara- 
tion of,  99 ;  ion,  148  ;  occurrence  of,  35  ; 
solubility  of,  in  water,  123 ;  uses  of,  44 

Hydrogen  bromide,  261  ;  constant-boiling 
solution  of,  263 

Hydrogen  chloride,  253 ;  composition  of, 
256;  constant-boiling  solution  of,  126, 
256 ;  oxidation  of,  249 ;  solubility  of,  in 
water,  123 

Hydrogen  cyanide,  291 

Hydrogen  fluoride,  246 

Hydrogen  iodide,  266 ;  constant-boiling 
solution  of,  267 ;  heat  of  formation  of, 
93,  334 ;  oxidation  of,  266 

Hydrogen  nitrate,  181 

Hydrogen  peroxide,  69;  catalysis  of,  71 ; 
equations  of  preparation  of,  100;  in  air, 
115 

Hydrogen  persulfide,  211 

Hydrogen  phosphides,  357 

Hydrogen  selenate,  231 

Hydrogen  selenide,  231 

Hydrogen  sulfate,  221 

Hydrogen  sulfide,  207 ;  action  of,  on  io- 
dine, 266;  action  of,  on  metals,  209; 
action  of,  on  oxygen,  209 ;  action  of,  on 
sulfur  dioxide,  213 ;  decomposition  of, 
by  heat,  209 ;  reducing  action  of,  209 ; 
solubility  of,  123 


,  INDEX 


549 


Hydrogen  tellurate,  232 
Hydrogen  telluride,  232 
Hydrogen  tellurjte,  232 
Hydrolysis,  226 ;  definition  of,  63 ;  of  acid 

salts,  227;  of  ferric  chloride,  470;  of 

normal  salts,  224 
Hydronitiic  acid,  175 
Hydrosol,  348 
Hydrosulfuric  acid,  209 
Hydroxides,  150,  377 ;  chief  reactions  of, 

377 ;  classes  of,  379 ;  solubility  of,  390 
Hydroxylamine,  184 
Hypobromites,  273 
Hypobromous  acid,  283 
Hypochlorites,  269;    preparation   of,  by 

electrolysis,  271 
Hypochlorous  acid,  252,  269;  anhydride 

of,  268 

Hypoiodites,  274 
Hyponitrous  acid,  184 
Hypophosphorous  acid,  363 
Hyposulfurous  acid,  227 
Hypothesis,  of  Avogadro,  310;  of  Prout, 

234 

Ice,  79;  manufacture  of  artificial,  174 

Iceland  spar,  420 

Ilmenite,  350 

Incombustible  substances,  26 

Indelible  ink,  498 

Indicators,  137 ;  table  of,  150 

Indigo,  297 

Indium,  450 

Inflammable  air,  35 

Infusible  white  precipitate,  493 

Inks,  467 

lodates,  274 

lodic  acid,  273 

Iodides,  267 

Iodine,  263  ;  action  of,  on  sodium  thiosul- 
fate,  400;  preparation  of,  from  kelp, 
264;  purification  of,  265;  tincture  of, 
265;  uses  of,  266 

Iodine  pentoxide,  273 

lodoform,  298 

Ionium,  533 

lonization,  141 ;  effect  of  dilution  on,  198 ; 
of  normal  acids,  155 ;  of  norm'al  bases, 
155 ;  of  normal  salts,  155 ;  of  polybasic 
acids,  150;  quantitative  measurement 
of,  146 ;  theory  of,  141 

Ions,  140;  electrically  charged,  142; 
sources  of  charge  upon,  143 

Iridium,  542 

Iridosmine,  542 

Iron,  459 ;  action  of,  on  water,  36 ;  action 
o£,  with  sulfur,  8  ;  cast,  460  ;  compounds 
df,  466;  see  also  Ferrous  and  Ferric 
compounds  ;  of  commerce,  459 ;  passive 
state  of,  465 ;  pig,  461 ;  pure,  459  ;  re- 
lation of  varieties  of,  464;  rusting  of, 
466 ;  tincture  of,  469 ;  wrought,  461 

Iron  carbonyls,  476 


Iron  cyanides,  472 

Iron  disulfide,  468 

Iron  family,  458 

Isomeric  compounds,  292 

Isomers,  292 

Isomorphous  substances,  82 

Kaolinite,  447 
Kelp,  iodine  in,  264 
Kerosene,  294 
Kilogram,  2 
Kinetic  theory,  51 
Kipp  apparatus,  38 
Krypton,  110 

Lactic  acid,  306 

Lactose,  299 

Lakes,  509 

Lampblack,  290 

Lanthanum,  451 

Latent  image,  497 

Laughing  gas,  186 

Law,  50;  of  Boyle,  46;  of  Charles,  48; 
of  chemical  combination,  83 ;  of  com- 
bining weights,  86 ;  of  conservation  of 
energy,  4 ;  of  conservation  of  mass,  83 ; 
of  Dalton,  125 ;  of  definite  composition, 
84;  of  Dulong  and  Petit,  319  ;  of  Gay- 
Lussac,  48,  310;  of  Graham,  50;  of 
heat  summation,  335  ;  of  Hess,  335  ;  of 
mass  action,  193 ;  of  multiple  propor- 
tion, 85  ;  of  osmotic  pressure,  135  ;  of 
Kaoult,  133,  310;  of  thermoueutrality, 
139 ;  periodic,  236 

Lead,  504  ;  acids  of,  510  ;  alloys  of,  375  ; 
compounds  of,  506;  red,  510;  sugar 
of,  508  ;  white,  508 

Lead  acetate,  508 

Lead  arsenate,  510 

Lead  basic  carbonate,  508 

Lead  bromide,  508 

Lead  carbonate,  508 

Lead  chloride,  508 

Lead  chromate,  510 

Lead  dioxide,  507 

Lead  iodide,  508 

Lead  metaplumbate,  510 

Lead  monoxide,  507 

Lead  nitrate,  507 

Lead  orthoplumbate,  510 

Lead  oxides,  506 

Lead  suboxide,  506 

Lead  sulfate,  509 

Lead  sulfide,  508 

Lead  tetrachloride,  508 

Leblanc  process,  250,  400 

Le  Chatelier,  principle  of,  202,  340 

Lepidolite,  392 

Levulose,  300 

Lime,  417 ;  air-slaked,  416  ;  slaked,  417 

Limekiln,  417 

Limestone,  419 

Lime-sulfur  spray,  211 


550 


GENERAL  CHEMISTRY 


Limewater,  417 
Limonite,  459 
Linde  machine,  78 
Liquefaction  of  gases,  77 
Liquids,  73 ;  freely  miscible,  126 ;  spar- 
ingly miscible,  128 
Litharge,  507 

Lithium,  392 ;  compounds  of,  393 
Lithium  bromide,  393 
Lithium  carbonate,  393 
Lithium  chloride,  393 
Lithium  phosphate,  393 
Lithophone,  509 
Litmus,  150 
Lunar  caustic,  498 
Lutecium,  451 

Magnalium,  431,  442 

Magnesia,  alba,  432  ;  usta,  431 

Magnesite,  430 

Magnesium,  430  ;  compounds  of,  431 

Magnesium  ammonium  arsenate,  369 

Magnesium  ammonium  phosphate,  433 

Magnesium  carbonates,  432 

Magnesium  chloride,  432 

Magnesium  family,  429 

Magnesium  hydroxide,  431 

Magnesium  nitride,  433 

Magnesium  oxide,  431 

Magnesium  phosphates,  433 

Magnesium  pyrophosphate,  434 

Magnesium  sulfate,  432 

Magnesium  sulfide,  433 

Magnetite,  459  ;  structure  of,  165 

Malachite,  484 

Malic  acid,  306 

Malt,  299 

Maltose,  299 

Manganates,  517 

Manganese,  514 ;  acids  of,  and  their  salts, 

517,  518 

Manganese  bronze,  515 
Manganese  compounds,  515 
Manganese  dioxide,  516 
Manganese  tetrachloride,  517 
Manganic  chloride,  516- 
Manganic  compounds,  516 
Manganic  sulfate,  516 
Manganin,  515 
Manganites,  516 
Manganous  carbonate,  516 
Manganous  chloride,  516 
Manganous  compounds,  515 
Manganous  hydroxide,  516 
Manganous  oxide,  515 
Manganous  sulfate,  516 
Marcasite,  468 
Marl,  419 
Marsh  gas,  295 
Marsh's  test,  366 

Mass,  2  ;  law  of  conservation  of,  2,  83 
Mass  action,  and  strong  electrolytes,  199  ; 

law  of,  193 


Matches,  356 

Matter,  2 ;  three  states  of,  73  :  varieties 

of,7x 

Meerschaum,  430 
Melting  point,  79 ;  of  amorphous  bodies, 

80 

Mercuric  chloride,  491 
Mercuric  compounds,  490 
Mercuric  cyanide,  291,  492 
Mercuric  fulminate,  492 
Mercuric  iodide,  491 
Mercuric  nitrate,  492 
Mercuric  oxide,  17,  490 
Mercuric  sulfate,  492 
Mercuric  sulfide,  491 
Mercurous  bromide,  489 
Mercurous  chloride,  489 
Mercurous  compounds,  488 
Mercurous  hydroxide,  489 
Mercurous  iodide,  489 
Mercurous  nitrate,  489 
Mercurous  oxide,  489 
Mercurous  sulfate,  489 
Mercurous  sulfide,  489 
Mercurous    salts,  molecular    weight    of, 

488 
Mercury,  486  ;  ammonium  compounds  of, 

492  ;  freezing  of,  283 ;  purification  of, 

486 

-  Metallo-acid  elements,  525 
Metalloids,  233 
Metallurgy,  383;   electrical  methods  of, 

385 
Metals,  233,  382  ;  extraction  of,  from  ores, 

383 

Metaphosphates,  362 
Metaphosphates  as  fluxes,  362 
Metaphosphoric  acid,  362 
Metasilicates,  345 
Metasilicic  acid,  345 
Metastannates,  504 
Metastannic  acid,  504 
Meteorites,  458 
Methane,  295  ;  heat  of  formation  of,  337  ; 

structure  of,  293 
Methyl  orange,  150 
Mica,  347 
Microcline,  447 
Microcosmic  salt,  361 
Milk,  299 

Milk  of  sulfur,  206 
Milk  sugar,  299 
Millon's  base,  493 
Minerals,  383 
Minium,  510 
Mischmetall,  452 
Mispickel,  364 

Mixed  gases,  solubility  of,  125 
Mixed    liquids,    boiling    point    of,    127 ; 

vapor  pressure  of,  126 
Mixed  metal,  452 
Mixtures,  13  ;  explosive,  329 
Mohr's  salt,  468 


INDEX 


551 


Molar  solutions,  128 

Molecular  weights,  309 ;  boiling-point 
method  for  determining,  311  ;  experi- 
mental calculation  of,  313  ;  freezing- 
point  method  for  determining,  314 ; 
method  of  Dumas  for  determining,  313 ; 
method  of  Victor  Meyer  for  determin- 
ing, 313 ;  of  elements,  318 ;  standard 
for,  311 

Molecule,  52,  96 

Molybdenite,  529 

Molybdenum,  528 

Molybdic  acid,  529 

Molybdic  oxide,  529 

Monazite,  451,  512 

Mordant,  444 

Mortar,  418 

Mosaic  gold,  503 

Moth  balls,  297 

Multiple  proportion,  law  of,  85 

Naphtha,  294 

Naphthalene,  297 

Nascent  state,  182 

Native  occurrence  of  elements,  383 

Neodymium,  451 

Neon,  111 

Neoytterbium,  451 

Nernst  lamp,  452 

Nessler's  precipitate,  493 

Nessler's  reagent,  491 

Neutralization,  148 ;  a  definite  act,  153  ; 
definition  of,  152  ;  heat  of,  157  ;  thermo- 
chemistry of,  157 

Newton's  metal,  375 

Nickel,  475  ;  compounds  of,  476 

Nickel  carbonyl,  476 

Nickel  chloride,   476 

Nickel  nitrate,   476 

Nickel  sulfate,  476 

Nickel  sulfide,  476 

Niobium,  525 

Niton,  533 

Nitrates,  182 ;  solubility  of,  390 

Nitric  acid,  175 ;  chemical  conduct  of, 
179  ;  constant-boiling  solution  of,  128  ; 
formation  of,  from  air,  177  ;  salts  of, 
182 

Nitric  oxide,  186 

Nitrifying  bacteria,  107 

Nitrites,  184 

Nitrobenzene,  297 

Nitrocellulose,  300 

Nitrogen,  104  ;  assimilation  of,  by  plants, 
107  ;  chemical  conduct  of,  106 ;  com- 
pounds of,  107,  167  ;  compounds  of, 
with  hydrogen,  168  ;  critical  constants 
of,  76  ;  cycle  of,  in  nature,  189 ;  per- 
centage of,  in  air,  117  ;  solubility  of, 
123  ;  unstable  compounds  of,  167 ;  utili- 
zation of  atmospheric,  424 

Nitrogen  dioxide,  187 

Nitrogen  hexoxide,  185 


Nitrogen  pentoxide,  188 

Nitrogen  tetroxide,  187 

Nitrogen  trichloride,  251 

Nitrogen  trioxide,  188 

Nitroglycerin,  303  ;  explosion  of,  331 

Nitrous  acid,  183  ;  salts  of,  184 

Nitrous  oxide,  185 

Nitrosyl  chloride,  258 

Nitrosyl  sulfuric  acid,  219 

Nonmetals,  233 

Normal  acids,  377 

Normal  salts,  156  ;  action  of,  on  water,  224 

Normal  solutions,  154 

Oil  of  mirbane,  297 

Oil  of  vitriol,  221 

Oleic  acid,  306 

Olein,  307 

Oleomargarine,  307 

Olivine,  347 

Opal,  344 

Open-hearth  furnace,  463 

Open-hearth  process,  462,  463 

Ores,  383 

Organic  chemistry,  281 

Organic  matter,  decomposition  of,  189 

Orpiment,  369 

Orthoclase,  347,  447 

Orthophosphates,  361 

Orthophosphoric  acid,  361 

Orthophosphorous  acid,  363 

Orthosilicates,  345 

Orthosilicic  acid,  345 

Osmic  acid,  542 

Osmium,  541 

Osmotic  pressure,  133  ;  laws  of,  135  ;  qual- 
itative demonstration  of,  134 ;  quantita- 
tive measurement  of,  134 

Oxalic  acid,  306  ;  decomposition  of,  287 

Oxidation,  22  ;  and  reduction,  472  ;  an 
increase  in  valence,  471;  ionic,  471; 
products  of,  23  ;  slow,  26 ;  speed  of, 
27  ;  weight  relations  in,  24 

Oxides,  23  ;  action  of  water  on,  61  ;  oxi- 
dizing and  reducing  properties  of,  377  ; 
reduction  of,  by  aluminium,  384  ;  re- 
duction of,  by  carbon,  384  ;  reduction 
of,  by  hydrogen,  197 

Oxidizing  agent,  22,  371 

Oxyacetylene  blowpipe,  297 

Oxygen,  16 ;  chemical  conduct  of,  21 ; 
combination  of,  with  hydrogen,  40 ; 
critical  constants  of,  76;  equations  of 
preparation  of,  98;  history  of,  16; 
laboratory  preparation  of,  19 ;  per- 
centage of,  in  air,  116;  properties  of, 
21  ;  solubility  of,  123  ;  the  standard  for 
combining  weight,  89 

Oxyhydrogen  blowpipe,  41 

Ozone,  31 ;  formation  of,  in  oxidation  of 
phosphorus,  32,  356;  relation  of,  to 
oxygen,  33 

Ozonizer,  31 


552 


GENERAL   CHEMISTRY 


Paint,  509  ;  luminous,  418 

Palladium,  541 ;  occlusion  of  hydrogen  by, 

39,  541 ;  use  of,  in  synthesis  of  water, 

67 

Palmitic  acid,  305 
Palmitin,  307 
Paper,  300 ;  manufacture  of,  216  ;  sizing 

of,  445 
Paraffin,  -294 
Paris  green,  368 
Parkes  process,  505 
Partial  pressure,  123 
Pearls,  419 
Pentane,  294 
Pentathionic  acid,  227 
Perchlorates,  272 
Perchloric  acid,  272 
Perchloric  anhydride,  268 
Perchlorides,  248 
Periodates,  274 
Periodic  acid,  274 
Periodic  families,  236 
Periodic  grouping,  234 
Periodic  law,  236  ;  value  of,  241 
Periodic  table,  237  ;  irregularities  in,  239 
Permanganates,  517 
Permanganic  acid,  517 
Peroxides,  507 
Persulfates,  229 
Persulfides,  211 
Persulfuric  acid,  228 
Petroleum,  294  ;  products  obtained  from, 

294 

Pewter,  375 
Phenolphthalein,  150 
Phlogiston  theory,  29 
Phosgene,  289 
Phosphates,  normal,  361  ;  primary,  361  ; 

secondary,  361  ;  solubility  of,  390  ;  ter- 
tiary, 361 
Phosphides,  357 
Phosphines,  357 
Phosphomolybdic  acid,  529 
Phosphonium  iodide,  358 
Phosphonium  salts,  358 
Phosphoric  acids,  361 
Phosphorous  acids,  363 
Phosphorus,  354 ;  acids  of,  361 ;  action  of, 

on  bases,  380  ;  action  of,  on  bromine, 

262;    metallic,   356;    oxides  of,    360; 

phosphorescence    of,    356 ;    red,    356 ; 

slow  combustion  of,  356  ;  yellow,  355  ; 

white,  355 

Phosphorus  halides,  359 
Phosphorus  pentachloride,  359 
Phosphorus  pentoxide,  360 
Phosphorus  sesquisulfide,  357 
Phosphorus  tetroxide,  360 
Phosphorus  trichloride,  359 
Phosphorus  trioxide,  360 
Phosphotungstic  acid,  531 
Photography,  497 
Pink  salt,  503 


Pitchblende,  531 

Plants,  effect  of,  on  composition  of  air, 
119 

Plaster  of  Paris,  421 

Platinates,  539 

Platinum,  539;  as  catalyzer,  216,  218; 
colloidal,  132 

Platinum  compounds,  540 

Plucker  tube,  112 

Plumbates,  510 

Plumbic  acid,  510 

Plumbites,  510 

Plumbous  acid,  510 

Pollucite,  410 

Polonium,  533 

Polyacids,  346 

Polysilicic  acid,  345 

Polymers,  187 

Porcelain,  456 

Potassium,  404 ;  compounds  of,  405 ; 
radioactivity  of,  535 

Potassium  bicarbonate,  408 

Potassium  bromate,  407 

Potassium  bromide,  406 

Potassium  carbonate,  407 

Potassium  chlorate,  19,  406 

Potassium  chloride,  406 

Potassium  chloroplatinate,  409 

Potassium  chromate,  523 

Potassium  cyanate,  409 

Potassium  cyanide,  409 

Potassium  dichromate,  523 

Potassium  ferricyanide,  472 

Potassium  ferrocyanide,  472 

Potassium  fluosilicate,  409 

Potassium  hydrogen  fluoride,  247 

Potassium  hydroxide,  207,  405 

Potassium  iodate,  407 

Potassium  iodide,  406 

Potassium  nitrate,  408 

Potassium  perchlorate,  409 

Potassium  permanganate,  518 

Potassium  sodium  cobaltinitrite,  409 

Potassium  sulfate,  407 

Potassium  sulfocyanate,  409 

Pottery,  456 ;  body  of,  457  ;  decoration  of, 
457 ;  glaze  of,  457 

Praseodymium,  457 

Precipitate,  130 

Precipitation,  and  equilibrium,  200  ;  the- 
ory of,  388 

Pressure,  critical,  76  ;  effect  of,  on  solu- 
bility of  gases,  124  ;  partial,  of  gases, 
123;  standard,  47 

Principle  of  Le  Chatelier,  202 

Properties  of  substances,  3 

Proteins,  307 

Prussian  blue,  473 

Prussic  acid,  291 

Puddling  furnace,  461 

Purple  of  Cassius,  539 

Pyrites,  468 

Pyrolusite,  516 


INDEX 


•  553 


Pyrophosphoric  acid,  362 
Pyrosulf  uric  acid,  228 
Pyrrhotite,  468 

Quartz,  344 
Quicklime,  416 
Quicksilver,  486 

Radicals,  149  ;  valence  of,  165 
Radium,  532  ;  disintegration  of,  532  ;  ef- 
fects of,  534;  emanation  of,  533;  energy 

of,  534 
Rare  earths,  451 ;   place  of,  in  periodic 

table,  240,  452 
Reactions,  8  ;   conditions  for  completing, 

199 ;  irreversible,  194 ;  speed  of,  191 ; 

types   of,    100;    volume    changes    in, 

320 

Realgar,  369 
Red  lead,  510 
Reducing  agent,  43 
Reduction,  43 ;  a  decrease  in  valence,  471 ; 

ionic,  471 
Rennin,  299 
Respiration,  118 
Retort  carbon,  323 
Reversion  of  phosphates,  364 
Rhodium,  541 
Rochelle  salts,  306 
Rose's  metal,  375 
Rubidium,  409 
Ruby,  443 
Ruthenium,  541 
Rutile,  350 

Safety  lamp,  330    • 

Sal  ammoniac,  411 

Sal  soda,  402 

Salt,  397 

Saltpeter,  408 

Salts,  137;  acid,  156;  basic,  156;  com- 
plex, 448  ;  definition  of,  157  ;  double, 
448 ;  general  properties  of,  152 ;  hy- 
drogen, 152  ;  insoluble,  390  ;  mixed, 
156,  449 ;  nomenclature  of,  159 ;  nor- 
mal, 156  ;  preparation  of,  from  oxides, 
153  ;  relative  ionization  of,  155 

Samarium,  451 

Saponification,  307 

Sapphire,  443 

Satin  spar,  421 

Saturated  compounds,  293 

Samarskite,  527 

Scandium,  451 

Scheele's  green,  368 

Scheelite,  529 

Schlippe's  salt,  373 

Schonite,  407 

Schweinf  urt  green,  368 

Selenic  acid,  231 

Selenium,  230 ;  in  glass,  455 ;  varieties  of, 
231 

Selenium  dioxide,  231 


Semipermeable  membrane,  134 

Separatory  funnel,  70 

Serpentine,  347,  430 

Siderite,  469 

Silica,  344 

Silicate  industries,  349,  453 

Silicates,  346;  fusion  of,  347;  solubility 

of,  390 ;  varieties  of,  347 
Silicic  acid,  347 ;  colloidal,  347 
Silicides,  342 
Silicon,  341 ;   acids  of,  345,  compounds 

of,  342 

Silicon  carbide,  348 
Silicon  dioxide,  344 
Silicon  fluoride,  343 
Silicon  halides,  342 
Silicon  hydrides,  342 
Silver,  494  ;  compounds  of,  496  ;  German, 

480 

Silver  acetate,  498 
Silver  arsenate,  369 
Silver  bromide,  497 
Silver  carbonate,  499 
Silver  chloride,  201,  497 
Silver  chromate,  499 
Silver  cyanide,  499 
Silver  fluoride,  496 
Silver  iodide,  497 
Silver  nitrate,  498 
Silver  nitrite,  498 
Silver  oxide,  496 
Silver  peroxide,  496 
Silver  phosphate,  499 
Silver  subchloride,  496 
Silver  suboxide,  496 
Silver  sulf  ate,  498 
.  Silver  sulfide,  499 
Slag,  384 
Smalt,  474 
Smithsonite,  434 
Soaps,  307 
Soda  ash,  400 
Soda  water,  284 
Sodium,  393;    action  of,  on  water,  36; 

compounds  of,  395 
•  Sodium  acetate,  437 
Sodium  bicarbonate,  402 
Sodium  bromide,  398 
Sodium  carbonate,  400 
Sodium  chloride,  397 
Sodium  chromate,  523 
Sodium  cyanide,  403 
Sodium  dichromate,  523 
Sodium  hydride,  395 
Sodium  hydrogen  carbonate,  402 
Sodium  hydroxide,  396 
Sodium  hyposulfite,  399 
Sodium  iodate  in  caliche,  264 
Sodium  iodide,  398 
Sodium  nitrate,  403 
Sodium  oxide,  395 
Sodium  peroxide,  20,  395 
Sodium  phosphate,  403 


554 


GENERAL  CHEMISTRY 


Sodium  polysulfide,  398 

Sodium  pyroantimonate,  404 

Sodium  sulfides,  398 

Sodium  sulfites,  399 

Sodium  thioantimonate,  373 

Sodium  thioarsenate,  370 

Sodium  thiosulf  ate,  399 

Sodium  tungstate,  530 

Sodium  vanadate,  527 

Solid  bodies,  79 

Solid,  definition  of,  73 ;  solution  of,  131 ; 
solution  of,  in  liquids,  128,  130;  vapor 
pressure  of,  80 

Solubility,  130 ;  effect  of  temperature  on, 
130 ;  of  liquids  in  liquids,  216 ;  of  mixed 
gases,  125 ;  of  very  soluble  gases,  124. 
See  also  Solubility  of  gases 

Solubility  curves,  130 ;  breaks  in,  131 ;  of 
gases,  124 

Solubility  of  gases,  123 ;  conditions  affect- 
ing, 123;  effect  of  pressure  on,  124; 
effect  of  specific  properties  on,  124; 
effect  of  temperature  on,  124 

Solubility  product,  388 

Solution,  and  speed  of  reaction,  193 ;  heat 
of,  338 ;  rate  of,  129 

Solutions,  122  ;  colloidal,  132 ;  color  of, 
due  to  ions,  146 ;  constant-boiling,  127  ; 
freezing  point  of,  133;  molar,  128; 
normal,  154  ;  of  constant-boiling  point, 
126 ;  of  electrolytes,  137 ;  of  gases  in 
gases,  122 ;  of  gases  in  liquids,  123 ;  of 
solids  in  liquids,  128 ;  of  solids  in  solids, 
131 ;  saturated,  128 ;  supersaturated, 
129  ;  theory  of,  136 

Solvay  process,  401 

Specific  heat  and  atomic  weight,  319 

Spectroscope,  112 

Spectrum,  111 ;  absorption,  113;  continu- 
ous, 111 ;  emission,  112 ;  vacuum-tube, 
112 

Spectroscope,  methods  of  use  of,  112 

Speed  of  oxidation,  27 ;  of  reaction,- 191 

Spelter,  424 

Sperrylite,  539 

Sphalerite,  434,  436 

Spiegel  iron,  515 

Spinel,  443 

Spinel  minerals,  444 

Spirits  of  hartshorn,  168 

Spodumene,  392 

Spray,  lime-sulfur,  211 

Stalactites  and  stalagmites,  420 

Stannates,  504 

Stannic  acid,  504 

Stannic  chloride,  503 

Stannic  compounds,  503 

Stannic  halides,  503 

Stannic  hydroxide,  503 

Stannic  oxide,  503 

Stannic  sulfide,  503 

Stannous  chloride,  502 

Stannous  compounds,  502 


Stannous  hydroxide,  502 

Stannous  oxide,  502 

Stannous  sulfide,  502 

Standard  cell,  490 

Standard  conditions,  49 ;  equation  for,  75 

Standard  pressure,  47 

Starch,  300 ;  color  with  iodine,  265 

Stassfurt  salts,  404 

Stearic  acid,  305 

Stearin,  307 

Steel,  462;  alloys,  465;  tempering  of, 
464 

Stibine,  371 

Storage  cell,  511 

Strontianite,  426 

Strontium,  426 

Strontium  bromide,  426 

Strontium  carbonate,  427 

Strontium  chloride,  426 

Strontium  compounds,  426 

Strontium  hydroxide,  426 

Strontium  iodide,  426 

Strontium  nitrate,  426 

Strontium  oxide,  426 

Strontium  sulfate,  426 

Strontium  sulfide,  426 

Sublimation,  88^* 

Substitution,  100 

Sucrose,  298 

Sugar,  298 ;  decomposition  of,  by  heat,  8 ; 
of  lead,  508 

Sulf ates,  224  ;  preparation  of,  224  ;  solu- 
bility of,  390 

Sulfides,  210 ;  reduction  of,  by  metals,  385 ; 
solubility  of,  390  ^ 

Sulfites,  215 ;  solubility  of,  390 

Sulfur,  203;  action  of,  on  bases,  380; 
action  of,  on  metals,  206;  action  of, 
on  nonmetals,  207  ;  action  of,  with  oxi- 
dizing agent,  207 ;  amorphous,  205 ; 
flowers  of,  204  ;  heat  of  combustion  of, 
333 ;  valence  of,  230 ;  varieties  of,  204 

Sulfur  dioxide,  92,  123,  212 

Sulfur  lac,  206 

Sulfur  monochloride,  229 

Sulfur  tetrachloride,  229 

Sulfur  trioxide,  216 

Sulfur  waters,  207 

Sulfuric  acid,  217;  action  of,  on  metals, 
222 ;  action  of,  on  organic  compounds, 
223 ;  action  of,  on  salts,  223 ;  dissocia- 
tion of,  by  heat,  221 ;  electrolysis  of, 
144;  fuming,  221;  hydrates  of,  223; 
oxidizing  properties 'of,  221;  salts  of, 
224  ;  structure  of,  223 

Sulfurous  acid,  213 

Sulfuryl  chloride,  229 

Sylvite,  406 

Symbol,  13,  90 

Symbol  weights,  90 

Talc,  430 
Tantalates,  527 


INDEX 


555 


Tantalum,  527 ;  compounds  of,  528 

Tantalum  lamp,  528 

Tartaric  acid,  306 

Telluric  acid,  232 

Tellurides,  231 

Tellurium,  231 

Tellurium  dioxide,  232 

Tellurium  trioxide,  232 

Temperature,  absolute  scale  of,  48 ;  effect 
of,  on  speed  of  reaction,  192;  criti- 
cal, 76 ;  effect  of,  on  speed  of  oxi- 
dation, 27 ;  limits  of,  reached  in  com- 
bustion, 339;  relation  of,  to  pressure, 
49 

Terbium,  451 

Tetrathionic  acid,  227 

Thallium,  450 

Theory,  51 ;  atomic,  93 ;  formation  of  a, 
51 ;  ionization,  141 ;  kinetic,  51 ;  value 
of  a,  52,  97 

Thermite,  442 

Thermochemistry,  332 

Thermoneutrality,  139 

Thiocarbonates,  290 

Thiocarbonic  acid,  290 

Thiocyanates,  292 

Thiocyanic  acid,  292 

Thionyl  chloride,  229 

Thiosulfuric  acid,  227 

Thiourea,  290 

Thorite,  512  -^ 

Thorium,  512;  radioactivity  of,  513, 
534 

Thorium  nitrate,  512 

Thulium,  451  * 

Tin,  500  ;  alloys  of,  375,  480 ;  compounds 
of,  501  (see  also  Stannous  and  Stannic 
compounds) 

Tin  plague,  501 

Tin  salt,  502 

Titanium,  350 ;  compounds  of,  350 

Toluene,  294 

Topaz,  oriental,  443 

Townsend  cell,  396 

Transformation  diagram,  80 

Transition  point,  81.  131 

Triads  of  Dobereiner,  234 

Tridymite,  344 

Trir  -ubstances,  82 

ads,  294 

Trisir  -ic  acid,  347 

I,  227 

29;  compounds  of,  530 
lamp,  530 
oil's  bh    .  423 
ntine,    action    of,    with    chlorine, 

T\i-e  elements,  238 
Type  metal    }7o 

•  ;ne.  448 
.  79 
iturated  compounds,  293 


Uranium,  531 ;  compounds  of,  531 ;  radio- 
activity of ,  531 ;  the  source  of  radium,  532 
Uranyl  salts,  531 
Urea,  289 

Valence,  161 ;  applications  of,  166 ;  de- 
termination of,  164;  nature  of,  165; 
relation  of,  to  acid  and  basic  conduct, 
378 ;  representation  of,  163 ;  unit  of, 
161 ;  variable,  162 

Vanadinite,  527 

Vanadium,  526 ;  acids  of,  527 ;  compounds 
of,  527 

Vanadium  alums,  527 

Vanadium  pentoxide,  527 

Vanadium  sulfate,  527 

Vapor  density,  313 

Vapor  pressure,  74  ;  correction  for,  in  gas 
measurements,  74;  determination  of, 
75 ;  lowering  of,  by  solute,  133 ;  of 
mixed  liquids,  126 ;  of  solids,  80 ;  values 
of,  Appendix  B 

Vapor-pressure  curves,  80 

Vaporization,  heat  of,  76 

Vaseline,  294 

Venetian  red,  469 

Vinegar,  305 

Vitrified  brick,  456 

Vitriol,  blue,  483  ;  green,  467  ; 

Vitriols,  467 

Volume,  calculations  of,  102      ^^^ 
perature  on,  320 ;  gram-moIe(S^f,  312 

Washing  soda,  402      * 

Water,  54 ;  action  of,  with  chlorine,  252  ; 
action  of,  on  elements,  61,  466 ;  action 
of,  on  oxides,  61 ;  analysis  of,  56,  199  ; 
catalytic  action  of,  64 ;  chemical  con- 
duct of,  60 ;  city  filtration  of,  58 ;  com- 
position of,  by  analysis,  64 ;  composition 
of,  by  synthesis,  65;  composition  of, 
by  volume,  68  ;  critical  constants  of, 
decomposition  of,  by  heat,  61,  339 ; 
decomposition  of,  by  iron,  36 ;  detection 
of  impurities  in,  56  ;  disease  transmitted 
by.  55 ;  effect  of  boiling  on,  58  ;  effect 
of  filtration  of,  58  ;  effect  of,  on  health, 
55 ;  electrolysis  of,  17;  exact  compo- 
sition of,  64  ;  freezing  point  of,  lowered 
by  electrolyte,  138  ;  heat  of  decompo- 
sition, 93  ;  heat  of  formation,  333  ;  heat 
of  fusion,  79 ;  heat  of  vaporization,  76 ; 
microorganisms  in,  55;  of  crystalliza- 
tion, 63,  99 ;  of  ocean,  55 ;  permanent 
hardness  of,  422  ;  properties  of,  60 ; 
purification  of,  444  ;  purification  of,  by 
distillation,  56 ;  sanitary  purification  of, 
58;  self -purification  of,  59;  softening 
of,  422 ;  temporary  hardness  of,  422 ; 
vapor  of,  in  air,  115,  118 

Water  glass,  346 

Weathering  of  rock,  119 

Weight,  2  ;  of  a  liter  of  a  gas,  321 


556                                 GENERAL  CHEMISTRY 

Weights,  atomic.  90  ;  atomic  and  molecu-  Yttrium,  451 

lar,  96  ;  formula,  91  ;.- gram-atomic,  97  ; 

gram-molecular, '97;  Symbol,  90  Zinc,  434;    compounds  of,  435;    white, 

Weston  cell,  490 '  435 

Whisky,  303  Zinc  carbonate,  436 

White  lead,  508  Zinc  chloride,  436 

Willemite,  347,  434                                     •  Zinc  dust,  434 

Wolframite,'  529  Zinc  hydroxide,  436 

Wollastonite,  347,  425  Zinc  oxide,  435 

Wood  distillation,  279  Zinc  sulfate,  436 

Wood's  metal,  375  Zinc  sultide,  436 

Wulfenite,  529  Zincite,  434 

Zircon,  351 

Xanthoprotein,  179  Zirconium,  351 

Xenon,  111  Zymase,  302 


APPENDIX  A 


LIST  OF  ELEMENTS,  THEIK  SYMBOLS  AND  ATOMIC  WEIGHTS 


ELEMENT                 SYMBOL 

ATOMIC 
WEIGHT 

ELEMENT 

SYMBOL 

ATOMIC 
WEIGHT 

Aluminium    .     .     . 

Al 

27.1 

Molybdenum  .     .         Mo 

96.0 

Antimony      ...         Sb 

120.2 

Neodymium    .     .         Xd 

144.3 

Arsron  .                     •         A 

39.88 

Neon      ....         Ne 

20.2 

Arsenic     ....         As 

74.96 

Nickel    ....  |      Ni 

58.68 

Barium     .... 

Ba 

137.37 

Niton     .... 

Nt 

222.4 

Bismuth    .... 

Bi 

208.0 

Nitrogen     . 

N 

14.01 

11.0 

Osmium 

Os 

190.9 

Bromine   .... 

Br 

79.92 

Oxygen  .... 

() 

16.00 

Cadmium       ...         Cd 

112.40 

Palladium  .     .     .   j      Pd 

106.7 

Caesium     ....         Cs  . 

132.81 

Phosphorus     .     .  :      P 

31.04 

Calcium    ....         Ca 

40.07 

Platinum    .     .     .         Pt 

195.2 

Carbon      ....         C 

12.00 

Potassium             .  !      K 

39.10 

Cerium      ....         Ce 

140.25 

Praseodymium     .         Pr 

140.6 

Chlorine  .... 

Cl 

35.46 

Radium       .     .     .         Ra 

226.4 

Chromium     .     .     . 

Cr 

52.0 

Rhodium               .  \      Rh 

102.9 

Cobalt       .     .     .     ..        Co 

58.97 

Rubidium              .  j      Rb 

85.45 

Columbian!  ...         Cb 

93.5 

Ruthenium      .     .         Ru 

101.7 

Copper      ....         Cu 

63.57 

Samarium  ...         Sa 

150,1 

Dysprosium  ...         Dy 

162.5 

Scandium   ...         Sc 

44.1 

Erbium     ....         Er 

167.7 

Selenium     . 

Se              79.2 

Europium.     .     .     .         Eu 

152.0 

Silicon   .     ... 

Si              28.3 

Fluorine   .     .     .     .   ,      F               19.0 

Silver     .... 

Ag          107.88 

Gadolinium   .     .     .         Gd 

157.3 

Sodium  .     .     . 

Na 

23.00 

Gallium    .... 

Ga 

69.9 

Strontium  . 

Sr 

87.63 

Germanium  .     .     . 

Ge 

72.5 

Sulfur    .... 

s 

32.07 

Glucinum      ...         Gl 

9.1 

Tantalum   . 

Ta 

181.5 

Gold     i      An 

197.2 

Tellurium  . 

Te 

127.5 

Helium      .... 

He 

3.99 

Terbium     .     .     . 

Tb 

159.2 

Molmiuui       .     .     . 

Ho 

163.5 

Thallium    ...         Tl 

204.0 

Hydrogen       .     .     . 

H 

1.008 

Thoriitln     .     .     .         Th 

232.4 

Indium      .... 

In 

114.8 

Thulium     ...         Tin 

168.5 

Iodine        .... 

I 

126.92 

Tin    .      .     .      .    *.   1      Sn 

119.0 

Iridium     .... 

Ir 

193.1 

Titanium              .  j      Ti 

48.1 

Iron 

Fe 

55.84 

Tungsten    ...         W 

184.0 

Krypton    .... 

Kr 

82.9 

Uranium     . 

r 

238.5 

Lanthanum   . 

La 

139.0 

Vanadium  . 

V 

51.0 

Lead     

Pb 

207.10 

Xenon    .... 

Xc 

130.2 

Lithium                    .   \      Li 

6.94 

Ytterbium             ~i 

Lutecium       ...         Lu 

174.0 

(Neoytterbium)  J  i 

172.0 

Magnesium    .     .     .         Mg 

24.32 

Yttrium      ...         Yt 

89.0 

Manganese    .     .     .         Mn 

54.93 

Zinc  

Zn. 

65.37 

Mercury    .... 

Hg 

200.6 

Zirconium  .     .     . 

Zr 

90.6 

APPENDIX  B 


Tension  of  Aqueous  Vapor  expressed  in  Millimeters  of  Mercury 


TEMPERATURE 

16  .... 

17  .    .    .    . 

18  .... 

19  .... 

20  . 


PRESSURE 
.    13.62 
.    14.4 
.    15.46 
.    16.45 
,    17.51 


TEM  i-  E  u  AT  r  RE  PKESSURE 

21 18.62 

22 19.79 

23 21.02 

24 22.32 

25     ..  .    23.69 


Weight  in  Grams  of  1  Liter  of  Various  Gases  measured  under  Standard  Conditions  ; 

also  the  Boiling  Point  of  Each  of  the  Gases  under  a  Pressure  of  1  Atmosphere 

(Barometric  Reading  =  760  mm.) 

WEIGHT 
OF  1  LITER 

Acetylene 1.1621 

Air 1.2928 

Ammonia 0.7708 

Argon 1.7809 

Carbon  dioxide   .    .    .  1,9768 

Carbon  monoxide    .    .  1.2504 

Chlorine 3.1674 

Fluorine 1.697 

Helium 0.1782 

Hydrogen 0,08987 

Hydrogen  chloride  .    .  1.6398 


BOILING 
POINT 

-83.8° 


-  33.5° 

-  186.0° 

-  78.2° 
-190.° 

-  33.6° 

-  187.° 

-  268.7° 

-  252.7° 

-  82.9° 


Hydrogen  fluoride 
Hydrogen  sulfide 
Krypton  .... 
Methane 

WEIGHT 
<»F  1  LITER 

..  .    0.893 
.    .    1.5392 
.    .    3.708 
0  7168 

ROILING 
POINT 

+  19  4° 
-  61.6° 
-  151.7° 

—  164  ° 

Neon    

.    .    0.9002 

—  239.° 

Nitric  oxide     .    . 
Nitrogen  .... 

.    .    1.3402 
.    .    1.2507 

-  153.° 
—  195.7° 

Nitrous  oxide  .    . 
Oxygen    .... 

.    .    1.9777 
.    .    1.4290 

-  89.8° 
—  182.9° 

Sulfur  dioxide     . 
Xenon      .... 

.    .    2.9266 
.    5.851 

-  10.1° 
-  109.° 

Densities  and  Melting  Points  of  Some  of  the  Elements 


MELTING 

DENSITY 

POINT 

Aluminium    .    . 

2.65 

658.5° 

Antimony      .    . 

6.62 

630.° 

Arsenic      .    .    . 

5.73 

sublimes 

Barium      .    .    . 

3.75 

850.° 

Bismuth     .    .    . 

9.80 

269.° 

Boron    .... 

2.5  (?) 

2000.°  to  2500. 

Bromine     .    .    . 

3.102 

-  7.3° 

Cadmium  .    .    . 

8.64 

321.° 

Calcium     .    .    . 

1.55 

780.° 

Carbon,  Diamond 

3.52 

Graphite 

2.30 

-4000.°(?> 

Chromium      .  .  . 

6.50 

148(.». 

Cobalt    .... 

8.6 

1490.° 

Copper  ... 

8.93 

1082.G 

Gallium     .     .    . 

5.95     - 

30.  23 

Gold 

19.32 

1062.4° 

Iodine   .... 

4.95 

113.° 

Iridium      .    .    . 

22.41 

2300.  > 

Iron  

7.86 

1505.3 

Lead  . 

11.37 

327.° 

Lithium     .    .    . 

0.534 

186.° 

Magnesium    .    . 

1.74 

C33.° 

Manganese    .    . 

7.39 

1207.° 

MELTING 

if 

DENSITY 

POINT 

Mercury  "... 

13.56 

-  38.8° 

Molybdenum  .    . 

8.6      above 

2550.° 

Nickel'    .... 

8.9 

1452.° 

Palladium    .    .    . 

11.4 

1549.° 

Phosphorus,  yellow  1.83 

44.1° 

Platinum      .    .    . 

21.50 

1755.° 

Potassium    .    .    . 

0.862 

62.5° 

Radium  .... 

(?) 

700.° 

Selenium     .    .    . 

4.8 

217.° 

Silicon     .... 

2.3 

1450.°(?) 

Silver  

10.50 

960.° 

Sodium    .... 

0.971 

97.° 

Strontium    .    .    . 

2.54 

900.° 

Sulfur,  monoclinic 

1.96 

119.25° 

Sulfur,  rhombic  . 

2.06 

114.5° 

Tantalum    .    .    . 

16.6 

2910.° 

Tellurium    .    .    . 

6.25 

450.° 

Thorium  .... 

11.3 

1690.° 

Tin  

7.29 

232.° 

Titanium      .    .    . 

3.54  1800°  to  1850.° 

Tungsten     .    .    . 

18.72 

3000.° 

Vanadium   .    .    . 

6.02 

1620.° 

Zinc    . 

7.10 

419.43 

14  DAY  USE 

RETURN  TO  DESK  FROM  WHICH  BORROWED 
LOAN  DEPT. 


This  book  is  due  on  the  last  date  stamped  below,  or 
on  the  date  to  which  renewed. 
Renewed  books  are  subject  to  immediate  recall. 

*  2 

24Apr'59DF 

fceC'D  LD 

APK  a  7  1959 

4Jan6lPMZ 

f^afi 



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